year 2 psci notes 2018 · molecules (simple and giant molecular structures) [giant is just fyi]...

Chinchilla Notes 2018

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Year 2 PSCI Notes 2018 FYE EDITION

Content Page No. Chemistry 2 Atoms and Molecules 2 Periodicity 6 Chemical Bonding 14 Chemical Changes 21

Physics 29 Sound 29 Electricity 34 Light 42 Measurement 47


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Chemistry: Atoms and Molecules What is an atom?

• The smallest particle of an element (avg. 10 × 10−10 m) • Retains all chemical properties of that element • Make up matter • Identical to other atoms of same element • Different to other atoms of different elements

How can atoms be represented in models? Ball and stick models

• Used to: o Study how atoms are combined in a substance o Explain the properties of different substances o Predict properties of unknown substances

• Ball represents atom, stick refers to force holding atoms together

• Differentiated by: Size and colour of balls

Circles

• Used to: o Represent atoms in scale to one

another (compare size) • Differentiated by: Size and colour of

circles Chemical Symbols

• Most effective as it can cover the wide range of elements (there are limited numbers of sizes and colours)

• Differentiated by: Chemical Symbol (colour and size if used with circles)


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How can atomic structure be represented in models?

• Atoms are made up of even smaller particles – sub-atomic particles

• There are three kinds – protons, electrons and neutrons

• Neutrons and protons have

the same relative mass

while electrons have

negligible mass → Mass of

nucleus contributes to the

mass of an atom

• All of these sub-atomic particles are held together by electrostatic

force of attraction

• Atoms contain an equal number of protons and neutrons, which

carry a +1 and -1 charge respectively → Positive and negative

charges cancel out

• Atoms are electrically neutral …… until ions come in later ……

• Atomic structure can be represented in the periodic table, showing

the mass number and atomic number

• Each element has a different atomic number • Mass number – atomic number = Number of neutrons • Atomic number = Number of electrons and protons EACH


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Electronic Structure

• There are some things to note when drawing the full electronic

structure of an atom:

o No. of e = no. of p

o The first electronic shell can take 2 electrons, and

subsequently 8 electrons

o Draw the electrons in orbit around the

nucleus in the following order: up,

down, left, right

• Parts needed:

o Element drawn

o Number of p, e, n

o Electronic configuration (2.8.8)

o Actual drawing

Isotopes

• When writing electronic structure of an element or isotope, follow

the standard nuclear notation:

• Isotopes are atoms of the same

element, with same no. of p and e

but different no. of n → still

electrically neutral

• They have same chemical properties (same no. of e) but different

physical properties (diff. no. of n)

• There are many isotopes of an element, and the actual mass

numbers in the PT is the proportion of these

different isotopes in nature

• Let’s explain this in simple terms: If Element

A has two types of isotopes, with Isotope Y

having 2n and Z having 3n……

o If there is 25% of Y and 75% of Z

naturally, then the mass number should be ___ (see below)


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Molecules

• A molecule is made up of two or more atoms chemically combined • There are two types: Molecules of Elements and Molecules of

Compounds Molecules of Elements

o Made up of a fixed number of one type of atom chemically combined

o Many non-metal elements exist as molecules of elements

o E.g. Hydrogen gas molecule Molecules of Elements

o Made up of a fixed number of different types of atoms chemically combined

o Fixed ratio of atoms regardless of state

o E.g. Water molecule (Made up of 1 oxygen and 2 hydrogen atoms)

• We use chemical formulae to represent molecules • The chemical formula tells us the types of atoms and the number of

each type of atom in one molecule of substance.


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Periodicity

What is the Periodic Table?

• A tabular arrangement of chemical elements which are organized

by:

o Atomic/proton numbers

o Electronic

Configuration

o Recurring Chemical

and Similar Physical

Properties

• The elements are organized

in groups and periods

o Groups are made up of elements with the same number of

valence/outermost electrons → similar chemical properties

o Periods are made up of elements with the same number of

electron shells (E.g. 2,1, 2,2, 2,3 etc)

• The elements in yellow are metals, blue are non-metals and green

are metalloids, which share properties of both metals and non-

metals

• Groups between II and III are transition metals

*Physical Properties: Appearance, state, density, colour, melting and

boiling point etc

**Chemical Properties: How elements react to one another


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General Trends

Atom/Proton Number (across period)

• Across Period → by 1

• Elements are arranged in increasing atomic number, which

increases by 1 from left to right across a period

Size

• The size of atoms is measured by the atomic radii,

the distance between the centre of the nucleus and

the outermost electron shell

Size (down group)

• Down Group → size

• Down the group, the number of electrons increase in number of

electron shells, increasing atomic radii (think of planets around the

sun in orbit – the more planets in orbit there are, the larger the

radius)

Size (across period)

• Across Period → size

• Across the period, the number of electron shells remains constant

• While both electrons and protons increase, the effect of the

increasing number of protons is greater than that of electrons as

electrons are being added to the same shell

• The resulting electrostatic force of attraction between nucleus and

electrons increases

• The nucleus attracts the electrons more strongly, pulling the atom’s

shell closer to the nucleus, decreasing atomic radii


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In summary:

What are the specific properties of different groups of elements?

Group I: Alkali Metals

Components: Lithium, sodium, potassium, rubidium, caesium and

francium

Physical Properties:

• Solid at room temperature

• Bright, silvery appearance

• Good conductors of heat and electricity

• Malleable – can be hammered without breaking

• Soft – can be cut with a knife

• Trend: Increase in density and atomic radii down the group


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Chemical Properties:

• Lose one electron to attain full shell structure – form positively

charged ions

• React violently with oxygen and water

o With oxygen, forms white oxides

o With water, forms alkaline hydroxides and liberates

(releases) hydrogen gas

• Burn with characteristic flame colours during chemical reactions

and can thus be differentiated

o E.g. Sodium → Yellow, Lithium → Red, Potassium → Lilac

• Trend: Reactivity increases down the group

o As no. of electron shells increases, electrostatic attraction

between positively charged nucleus and negatively-charged

electrons weakens

o Valence electrons held more loosely and ease of losing

valence electron increases

General formula for chemical reactions with:

Water:

Oxygen: *a or aqus means dissolved in water


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Flame tests: Doing francium would probably destroy everything

Applications: (Just memorise 2 or 3)

• Sodium – Sodium vapour lamps • Caesium – Atomic clocks

o Extremely accurate, used for GPS and synchronization of the internet

• Potassium and Sodium – Cleaning Agents • Potassium Nitrate – Fertilizer • Sodium Fluoride – Toothpaste


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Group VII (7): Halogens

Components: Fluorine, chlorine, bromine, iodine, astatine

Physical Properties:

• Change in colour down group (learn this!)

• Highly toxic and have choking vapours

o Iodine in particular is an extremely volatile solid that sublimes

into vapour readily on application of heat at room temperature

(turns immediately from solid to gas)

• Trend: Melting and boiling points increase down the group

o Strength of the van der Waals forces (electrostatic forces

between atoms) increases since the atoms have more

electrons as you descend the group

▪ More energy required to break this force

*Ignore astatine as it is radioactive


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Chemical Properties

• Gain one electron to gain full shell structure to form a negatively

charged ion

• Do not exist naturally in elemental form but as salts in the halide

ions (not very important) (𝐹−, 𝐶𝑙− etc.)

• When they are in elemental form, they are toxic and reactive

diatomic gases (𝐹2, 𝐶𝑙2 etc.)

• Trend: Decreases down the group

o As number of electron shells increases, electrostatic

attraction between positively charged nucleus and negatively

charged electrons weakens

o Electrons are more easily gained

Applications: (just memorise 2 or 3)

• Fluorine – Toothpaste, disinfectant

• Chlorine – Disinfectant, pesticides

• Bromine – Fire retardants, insecticides

• Iodine – Silver iodide in photographic film, rain-

making by seeding clouds (like what we did

when the haze got too bad)

Group 0: Noble/Inert Gases Components: Helium, neon, argon, krypton, xenon and radon Physical Properties:

• Odourless • Colourless • Monoatomic gases at room temperature (He, Ne etc.) • Low boiling point

Chemical Properties:

• Complete valence shell → relatively unreactive → little tendency to gain or lose electrons

• Does not tend to combine with other elements (also non-flammable)


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Applications: (Just memorise 2 or 3)

• Helium – Fills airships, weather balloons due to lower density than air and non-flammability

• Neon – Neon lights and signs • Argon – Increase light bulb’s

life as it prevents reaction of tungsten which would otherwise wear away when reacting with oxygen at high temperature

• Krypton – Filler in double glazing as it provides optimal insulation by moving slower than air molecules due to being heavier

Summary of trends for various groups: Group I:

• Increase in density and atomic radii down the group • Reactivity increases down the group

Group VII:

• Melting and boiling points increase down the group

• Decreases down the group


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Chemical Bonding Overview In this year’s syllabus:

Ionic Bonding Covalent Bonding Force Electrostatic forces of

attraction between oppositely charged ions (EFOA)

Electrostatic forces of attraction between positively charged nuclei (+) and shared pair of electrons (-)

Between Metals and Non-Metals Non-Metals E.g. NaCl, MgSO4, CuCl2, ZnO CH4, NH3, HCl, H2O Present Particles

+ and – charged ions (cations and anions respectively)

Molecules (Simple and Giant molecular structures) [Giant is just FYI]

Ionic Bonding What are Ions?

• Ions are a single atom or a group of atoms which has an electric charge due to the loss or gain of electrons

o When an atom loses an electron, it has more positive charges (protons) than negative charges (electrons) → becomes positively charged ion

o In the same way, when an atom gains an electron, it has more negative charges than positive charges → becomes negatively charged ion

• The purpose of gaining or losing these electrons is to gain stability, as by themselves, they are unstable and reactive

o By gaining or losing these electrons, they gain fully-filled valence electron shells (FFVES), similar to noble gases ▪ E.g. Sodium, which is in Group 1, has one valence

electron (2,8,1) and has to take the shortest path to gain FFVES

• It is easier to lose electron than gain 7, so sodium becomes a positively charged ion denoted by Na+ because there are more protons than electrons


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• Since Noble Gases (group 8) already have FFVES, they are not involved in ionic bonding. Since elements in group 4 have to gain or lose 4 electrons, which takes up too much energy, they do covalent bonding instead

o As such, only Groups 1, 2, 3, 5, 6, 7 undergo chemical bonding • There are some exceptions, with some transition metals like iron

and copper forming charges (you don’t need to understand why yet) o The charges of these transition metals can vary, and are

denoted by roman numerals ▪ E.g. Iron (II) forms Fe2+ and iron (III) forms Fe3+

• There are two main groups of ions: monoatomic and polyatomic o Monoatomic ions are made of just one atom

▪ E.g. Sodium ion Na+ o Polyatomic ions are made of multiple atoms, often covalently

bonded, with a collective charge ▪ E.g. Ammonium ion NH4

+ • When two ions chemically combine, an ionic compound is formed,

which is neutral


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Here is the list of ions to memorise for exams (you may be asked to match these to form ionic compounds)

Positively-Charged Ions Negatively-Charged Ions

+

Sodium Na+ - Nitride N3-

Potassium K+ 2-

Oxide O2-

Silver Ag+ Sulfide S2-

Copper (I) Cu+

3-

Bromide Br-

Hydrogen H+ Chloride Cl-

+2

Barium Ba2+ Fluoride F-

Calcium Ca2+ Iodide I-

Copper (II) Cu2+

Polyatomic

Hydroxide ion OH-

Iron (II) Fe2+ Nitrate ion NO3-

Lead (II) Pb2+ Hydrogen sulfate or bisulfate

HSO4-

Magnesium Mg2+ Sulfate ion SO42-

Zinc Zn2+ Carbonate ion CO32-

Nitrite ion NO2-

+3

Aluminium Al3+ Sulfite ion SO32-

Iron (III) Fe3+ Hydrogencarbonate or Bicarbonate ion

HCO3-

Polyatomic Ammonium ion NH4+ Manganate (VII) ion MnO4

-

Dichromate (VI) ion Cr2O72-

Phosphate ion PO43-

Metals and Non-Metals Transition Metals Polyatomic Ions IP Only Misconception: “ide”, “ite” and “ate”

• At the end of certain ions, you might see “ide”, “ite” and “ate” at the end of the same element. These are only applicable for anions.

• As of now, you only need to know how this implies to nitrogen and sulfur

Nitrogen Sulfur Nitride N3- Sulfide S2-

Nitrite NO2- Sulfite SO3

2-

Nitrate NO3- Sulfate SO4

2-

As you can see:

• Ions that end with “ide” are the ions of an atom of a single element • Ions that end with “ite” are polyatomic ions with oxygen • Ions that end with “ate” are also polyatomic ions with oxygen, with one more

oxygen atom


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Examples of Dot-and-Cross Diagrams

• Dot and cross diagrams are common questions that you will find in papers, and is similar to drawing the full electronic structure of an atom, with the addition of:

o Square Brackets o Number of Ions o Charge Number and Sign

• These diagrams represent how different ions combine to form ionic compounds

1. Sodium Chloride

• This example is easier because each ion has an equal number of

positive and negative charge respectively (both have just one extra or missing electron) However, what happens when this number is unequal?


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2. Aluminium Oxide

Physical Properties of Ionic Compounds

• Binary ionic compounds (compounds made up of two elements) like NaCl do not exist in pairs, but form a giant crystal lattice structure

o The name “giant” does not refer to size but the millions of pairs of ions packed together

o Ions are closely packed, in a regular and repeating pattern and are held in place by ionic bonds

• The chemical formula of these structures can be deduced by telling the ratio of ions surrounding one another:


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• Here are the physical properties of these giant crystal lattice structures:

o High melting and boiling points ▪ A large amount of energy is required to overcome the

SEFOA between oppositely charged ions ▪ As a result, ionic compounds are solids at room

temperature and have very high melting and boiling points compared to covalent bonds

o Soluble in water but insoluble in organic solvent ▪ Water molecules are attracted to ions, causing ions to

be pulled from the lattice structure and weakening the electrostatic forces between ions → compound dissolves to form aqueous solution

▪ No water is present in organic solvents, causing ions to remain tightly held in the lattice structure, and ionic compounds to be insoluble

• Organic solvents are carbon-based, including methane and ethanol

o Can conduct electricity when in molten or aqueous solution ▪ In molten or aqueous form, ions are mobile and can

conduct electricity ▪ In solid form, there are no free-moving ions as ions are

held tightly in place in the lattice structure o Strong

▪ SEFOA between ions hold ions rigidly together, causing them to be strong (can withstand heavy loads)

o Brittle ▪ Brittleness refers to how easy it is to shift the position of

atoms or ions in a lattice ▪ Pressure on ionic compounds causes the layers to shift

slightly, causing ions of the same charge to come closer and the electrostatic repulsion between ions of same charge to break the structure

• This may seem contradictory, but take glass for example. It is strong and can withstand heavy loads (can thus be used in huge aquariums) but is easily broken when hit with a sudden force (e.g. hammer)


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Covalent Bonds

• Purpose of covalent bonding is also to achieve FFVES o This is instead done by sharing of valence electrons between non-

metal atoms (this only affects the valence shell) • Can take place between atoms of the same element

o Diatomic molecules (mainly from group 7), are formed by covalent bonds

o These diatomic molecules are hydrogen, oxygen, nitrogen, fluorine, chlorine, bromine, iodine and astatine (must memorise!!)

• There are several ways for covalent bonding to occur: o Covalent bonds can be single, double or even triple bonds (one

bond is signified by a pair of shared electrons – one from each involved atom) ▪ The greater the number of bonds, the greater the

electrostatic forces of attraction between atoms o The same element can make different types of bonds to form

different covalent compounds ▪ E.g. Oxygen can bind with another oxygen atom to form one

double bond oxygen gas O2 or bind with two hydrogen atoms to form two double bonds – water H2O

• By the end of bonding, every atom involved should have FFVES (use this to check your answer when drawing diagrams)

Examples

1. Fluorine Gas • Fluorine gas is an example of a diatomic molecule combining covalently


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2. Methane Gas

• However, in a more complex case like Methane Gas and limited time

during exams, there is a shortcut…… 😊

Physical Properties of Covalent Compounds

• Low Boiling and Melting Point o While atoms in the molecules are held together by strong

covalent bonds, the molecules are held together by weak intermolecular forces of attraction (van der Waal’s forces) ▪ Do note that when boiling or melting occurs, the

molecules just split from one another and not the atoms within the molecules, as this is a change of state and thus a physical change → chemical composition does not increase

o Suggested Answer: In a simple molecular structure, the small discrete molecules are held together by weak intermolecular forces of attraction. Little energy is required to overcome these forces.

• Insoluble in water, soluble in organic solvents (opposite of ionic) • Do not conduct electricity regardless of state

o No free moving ions or electrons to conduct electricity o They exist as molecules, neutral, no charge


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Chemical Changes Physical vs Chemical Changes

Physical Change Chemical Change No new substance formed New substance formed

→ Properties of products formed are completely different from that of the reactants

Reversible, temporary change Irreversible, permanent change E.g. Dissolving sugar in water, melting of ice cube, boiling of water, Clay moulded into a different shape (basically including every change of state)

E.g. Decomposition, curdling of milk, mixing acidic and basic components, frying an egg

Both may involve heat gained or released Both may have gas released or solid formed

Both may involve a change in colour In this chapter, the following chemical changes will be covered:

1. Interactions between matter and energy: i. Heat

ii. Oxygen iii. Light iv. Electricity

2. Interactions between matter (mixing): i. Acids + Alkali (Neutralisation)

ii. Acids + Metal iii. Acids + Carbonates iv. Metal + Water (Page 9) v. Metal + Oxygen (Page 9)

Interactions Between Matter and Energy

• When matter changes, there are also changes in energy o Energy is absorbed from the surroundings to the reactants to

break the bonds o Energy is released to the surroundings when new bonds are

Formed


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Heat

• The main reaction that takes place with heat is thermal decomposition

• A process in which a substance is broken down into two or more simpler substances by the effect of heat

• The name of the oxide is the same as that of the carbonate (calcium carbonate → calcium oxide)

• Apart from chemical changes, heat also causes physical changes like expansion and contraction and change in state

Oxygen

• Chemical changes caused by oxygen are called oxidation • There is no fixed formula for this, but here are some examples:


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Light

• Light is also quite straightforward and does not have a formula. Here is an example:

Electricity

• No formula, but allows for electroplating and electrolysis: o Electrolysis

▪ Electrolysis is the decomposition of substances (the electrolyte) by passing an electric current through it.

▪ An electrolyte is a liquid which can conduct electricity due to the presence of ions

o Electroplating ▪ Electrolysis can be used to coat an object with a metal, a

process called electroplating ▪ This is done by immersing an object in the electrolyte

and passing an electric current through it


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Interactions between matter Acids + Alkali (Neutralisation)

• Here is the formula and examples:

• Examples of Acids (acidic component caused by Hydrogen)

o Hydrochloric acid, HCl

o Sulfuric acid, H2SO4

o Nitric acid, HNO3 • Examples of Basics

(hydroxides) o Sodium hydroxide, NaOH o Potassium hydroxide, KOH o Aqueous ammonia/ammonium hydroxide, NH4OH

• Naming the Salt:

o The first part of the name comes from the metal / metal hydroxide / metal carbonate

o The second part of the name comes from the acid used ▪ Chloric → chloride ▪ Sulfuric → sulfate ▪ Nitric → nitrate

o As such: ▪ Hydrochloric acid + Sodium hydroxide → Sodium

chloride + water • Acid or Alkali?

o Universal indicator, indicated by litmus paper

o E.g. To test for acidity of gaseous hydrogen chloride, use a moist litmus paper or dissolve the gas in water and then test using litmus paper or any pH indicator


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Acid + Metal

• To test for hydrogen, cover the top of the test tube to prevent any hydrogen from escaping. Then, place a glowing splint in the test tube, which should be extinguished with a “pop” sound if hydrogen is present

Acid + Carbonate

• To test for carbon dioxide, bubble the gas into limewater (calcium hydroxide) and a white precipitate should be formed in the limewater


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Writing Chemical Equations

• Writing chemical equations requires the identification of the following:

o The chemical formula of the individual atoms involved or the compounds involved based on the chemical formula

o The state of each compound ▪ There are two sides of the equation: the reactants and

the products • Both sides must total up to have an equal number

of each element, as due to the Law of Conservation Mass, atoms can neither be created nor destroyed

▪ To figure out the reactants and products, try to match with the formulas above

• Here is an example of a chemical equation:

• Here are the steps to writing a chemical equation:


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Tips, Tricks and Rules of Writing Equations

1. Both sides must be equal 2. Individual elements within a compound cannot be multiplied by

themselves: a. E.g. Mg in MgO cannot be multiplied individually by two.

Instead, the whole MgO must be multiplied 3. Number of atoms per element is to be written in subscript – when a

compound or atom is multiplied, the number it is multiplied by must be placed in front of the formula:

a. E.g. If MgO is multiplied by 2, it becomes 2 MgO 4. General way to determine states:

a. Metals and Carbonates → solid b. Non-Metals → liquid and gas c. Acids, Alkali and Salts → aqueous

5. Use the chemical formula to help you: a. Di – two b. Tri - three c. Tetra – four d. Pent – five

i. E.g. dinitrogen pentoxide (N2O5) 6. Practise!!!

Note:

• In exam questions: o “Identify” – CO2, carbon dioxide o “Chemical name” – Carbon dioxide o “Chemical Formula” – CO2


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Physics: Sound Sound is produced by vibration and requires a medium to propagate. (Propagate is the transfer of sound) Longitudinal waves is a wave vibrating in the direction of propagation.

• The vibration of the air moves outward in all directions in a form of a

wave. In this way, the sound is propagated from the source to the

surrounding.

• Energy is transmitted when sound propagates.

• Sound is needed for communication in our daily lives

It travels through these medium at different speeds. (From Fastest to Slowest) When molecules at one end start to vibrate, energy is transferred to the neighboring molecules…

• Solid very rapidly. This is because the molecules are packed very

closely together. Speed of sound in solids = 5000-6000 m/s

• Liquid less rapidly. This is because the molecules of a liquid are not

packed as closely together as a solid.

Speed of sound in liquids = 1500 m/s

• Gas slowly. This occurs within a gas as the molecules are far

apart from one another. Speed of sound in air = 330 m/s

Similarly, in a vacuum pump with an alarm clock inside, the ringing of the bell goes silent. This shows that sound can travel only in matter. Sound cannot travel in a vacuum. The vibration of the air moves outward in all directions in a form of a wave.

• 1 Cycle is made up of two semi-circles forming one circle.

• A Compression is a region of a wave with particles closer together.

(Higher air pressure)

• A Rarefaction is a region of wave with particles further apart.

(Lower air pressure)


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• A Wavelength is the distance between two adjacent compressions

or rarefactions. (It is commonly designated in Lambda- λ )

(Wavelength) Amplitude/Loudness Amplitude affects the loudness of sounds. The Width is higher when it is louder and is shorter when it is softer (It is measured in Decibels- dB) Therefore,

• Louder = Higher Decibels

• Softer = Lower Decibels

• Prolonged 85-100dB can cause hearing loss.

• Safe sound before sustaining permanent damage – 80-95dB.

Frequency/Pitch Frequency -> pitch is the number of vibrations per Millisecond (It is measured in Hertz – Hz) . Therefore,

• High Frequency = Higher pitch (More cycles per Mil-Sec)

• Low Frequency = Lower pitch (Less cycles per Mil-Sec)

• One fact is we can typically hear and detect sound ranging from 20

Hz to 20,000 Hz (20 kHZ).


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The wave energy dissipates as the wave travels, causing the sound to decrease in volume. There are many Factors that affect the speed of the sound

• Temperature: The speed of sound is faster in a hotter area then a

colder area. This is due to how when the temperature is increases,

the particles in the medium gain more kinetic energy, vibrating and

colliding with neighboring particles faster. Therefore, with the

energy transferred faster, the speed of sound is increased.

• Humidity

• Type of medium

• Density

To measure sound directly, we use an oscilloscope (It takes a picture of the sound)

It can be measured by the formula: (2 x Distance)/Time taken

• The Echo Method, was first used by Marin Mersenne to measure the

speed of sound in air in 1600 by measuring the return of the echo.


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The Human Ear

1. The external ear is the cup-like flap that collects and channels the

sound to the middle ear. The outer ear also provides protection for

the middle ear.

2. A tube called the ear canal channels the sound to the eardrum. This

is a flap of skin that vibrates according to the frequency of the sound

coming in.

3. The vibrations are then transferred by three bones (hammer, anvil

and stirrup -respectively) to the inner ear

4. Here, the vibrations are transferred to the cochlea. The cochlea

begins to vibrate according to the different frequencies. The

cochlea’s inner surface is lined with nerve cells that transform

vibrations (Compression waves) to nerve impulses.

Ear drum – Vibrates when sound waves reach the ear Auditory nerve - carries electrical signals to the brain

• A Typical human can hear sound

from 20 Hz to 20 kHz (20,000 Hz).

Sound Higher than human audible

range is called ultrasound.

Sonar • Reflected sounds (echoes) from

the oceans floor can provide

information about underwater

conditions as well as objects

surrounding the ship. Ship use

sonar to detect shoals fish and

large whale.

• Sonar is a technique that used sound to gather information about

the environment

1. A ship emits a sound that travels down to the ocean floor

2. The time interval for the sound to go from the ship to the oceans

floor and back to the ship

3. 2 x distance travelled = speed of sound x time taken


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Ultrasound and technology

• Ultrasound is used in hospitals to observe the internal body without

the need for an operation

• Ultrasound is defined as sound with frequencies above 20,000 Hz,

which cannot be heard by the human ear

• For example, ultrasound allows doctors to observe foetuses in

pregnant mothers, detect cancers and observe tumours within the

body. It is also used to monitor blood flow in the veins and

abnormalities with the heart.

Doctors also use ultrasound to break up kidney stones into small pieces so that they can be passed out of the body easily, thus avoiding surgical operation.

Popular Exam questions:

1. Q- Why is the echo of smaller amplitude?

A- Sound energy has been lost to the surroundings during the

propagation of the sound.

2. Q- Why is a hard surface used as a sound reflector?

A- Hard surface reflector ensures all sound waves are reflected and not absorbed by the material.

3. Q- Suggest how temperature affects the speed of sound.

A- Temperature increases, the particles in the medium gain more kinetic energy move faster and collide with each other more, sound energy travels faster hence speed of sound is increased.

4. Q- Why can’t you hear sound in space?

A- Sound needs a medium to pass through. However, in space, there

is no matter/medium. Thus, no sound can be heard

5. Q- In Kungfu serial, the actors place their ears close to the ground to

listen to enemy on horseback approaching. Explain why?

A- By placing their ears nearer to the ground, he might be able to

hear the vibrations from the ground clearer as sound travels the

slowest through the air and fastest through solids.


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Electricity Flow of Electricity Electric current (I) For an electrical appliance to work, electricity must flow through it. The flow of electricity is called an electric circuit. An electric current is the rate of flow of electric charge.

• An electric current is measured by the amount of electricity charge

moving per unit time past any point in the circuit.

• The SI unit of current is the ampere (A). One ampere of current

means that one unit of charge flows in one second.

In a closed Circuit, the current flows from the positive terminal of a battery to the negative terminal. This is known as a the conventional current

After the discovery of electrons, it became known that electron flow occurs from the negative to the positive terminal.

• The direction of conventional current flow is opposite to the

direction of electron flow.

Electrical circuit • All the components of the electric circuit must be connected

correctly for it to work.

• A circuit diagram helps us see if the electrical components are

connected correctly.

Drawing circuit diagrams • We use common

symbols when

drawing circuit

diagrams.

Direction of electron flow

Direction of conventional current flow


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Potential difference (p.d.)/ Voltage (V) Electric charges need energy to push them along the circuit. Voltage is the amount of energy needed to push 1 coulomb of electricity.

• Water always flows from higher to lower ground. Similarly, a

positive charge always flows from a point of higher potential to a

point of lower potential. An electric current can flow only when

there is a potential difference (V) or p.d.

• The potential difference (or p.d. ) between any two points is the

amount of energy needed to move one unit of electric charge from

one point to the other.

• The SI unit of potential difference is a Volt (V)

• One volt of potential difference means that one joule of energy is

needed to move one unit of charge. The more energy needed to

move charge between two points in s circuit, the greater the

potential difference between the two points.

• V (Unit for p.d.) = J (Unit for energy) ÷ C (Unit for charge)

Resistance (R) An electrical component resists or hinders the flow of electric charges when it is connected in a circuit. In a circuit component, the resistance to the flow of charge is like how a narrow channel resists the flow of water. The resistance of a component is the ratio of the potential difference across it to the current flowing through it.

• The SI unit of resistance is the ohm (Ω)

• The greater the ratio of V to I, the greater is the resistance.

Factors that affect resistance: • Thickness (Cross-sectional area) - The thicker the wire/material,

the lower the resistance in the wire. With a thicker wire/material, it

is easier for the charges to flow through, because there are more

electrons to carry the current.

• Length - The longer the wire/material, the greater the amount of

resistance. The longer the wire/material, the electrons collide with

the ions more often.

• Material – The type of material affects the amount of resistance the

material has. The material that has better conductor of electricity

has a lower resistance. E.g. Copper/Iron has a lower resistance in

comparison with flexible plastic which has high resistance.


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Adding Resistors in Parallel decreases the resistance, thus the bulb would be brighter Adding Resistors in series increase the resistance, thus the bulb would be dimmer. Summary of Energy, Current, Voltage and Resistance:

Energy (J)

Current (I)

Voltage/Potential difference (p.d)

Resistance

SI Unit Jolts Ampere Volt Ohm Symbol for Unit J A V Ω Measuring Instrument

-Nil- Ammeter Voltmeter Fixed resistor

Symbol for Instrument

-Nil-

Instrument connection (Series/Parallel)

-Nil- Series Parallel Series and Parallel

V=IR Effects of an Electrical Current Magnetic effects Electromagnetic strength depends on:

• Number of coils around the nail per unit length. The greater the

number of coils around the nail per unit length, the stronger the

magnetic attraction of the nail.

• The Power Supply. By increasing the Power Supply, the stronger the

magnetic attraction of the nail.

Application: Telephones • In telephones, a changing magnetic effect causes a thin sheet of

metal (diaphragm) to vibrate.

• The diaphragm is made of a metal that can be attracted to magnets.

• The diaphragm becomes attracted to the electromagnets. As the

person on the other end of the line speaks, his voice causes the

current in the circuit to change. This causes the diaphragm in the

earpiece to vibrate, producing sound.


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Chemical effects Electrolysis of acidified water When an electric current is passed through certain liquid, a chemical change can be observed. This chemical effect is made use of in electrolysis of metals. Electricity can’t flow through pure water but can flow when minerals are mixed inside the Pure water. When electricity passes through acidified H2O (Water), the water particles would split up into H2 (Hydrogen) and O2 Oxygen.

• The battery acts like an ‘electron pump’. It draws electricity from

anode (+) to supply it to the cathode (-).

• Anode (+) – O2 (Oxygen)

• Cathode (-) – H2 (Hydrogen)

Electroplating Whatever you coat/plate, the liquid must contain the ion of that material.

• Many objects are coated with a metal by immersing the object in a

liquid and then passing an electric current through the liquid. This

method of coating objects with metals is called electroplating.

Increase the voltage/Current/emf or concentration of the solution to plate an object more quickly.


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Heating and lighting effects When an electrical current pass through a wire, the electrical energy is converted to heat. The heating elements are made of materials with high melting points (Iron, nickel and chromium). The heating effect of a current can lead to the lighting effect. The current that flows through the filament in the light bulb heats up. The heating is so hot, the filament glows and gives off light. Why must the filament of the bulb be coiled? The longer the filament, the more light it gives off. (more surface area for heating) Hazards of electricity/ Household Electricity Power The power of an electrical circuit is the amount of electrical energy converted to other forms of energy per unit time by the component. The SI unit of power is watt (W). One watt of power means in 1 second, 1 joule of electrical energy is converted to other forms of energy.

• Kilowatts (kW) – 1 kw = 1,000 W

• Megawatts (MW) – I MK = 1,000,000W

• Gigawatts (GW) – 1GK = 1,000,000,000W

• Energy= Power x Time

10W -> 10 Joules/ 1 sec Therefore, 10W= 10 J ÷ 1 Sec

Electrical appliances safely convert electrical energy into useful energy only if they Are not damaged. If we use damaged appliances, hazards such as electrical fires and electrocution may occur.

Electrical fires Electric shocks and electrocution • Irons and kettles require

large currents to produce heat. The wires in most electrical systems do not heat up as the components resist the flow of electric charge

• Large currents occur when electrical circuits are damaged, do not work

• Currents are dangerous when the pass through a person’s body. When this happens, he or she will experience an electric shock or electrocution. This usually results in serious injury or death.


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properly, or misused. For example, if a power outlet is overloaded, this may draw a large current. The large current generates heat, which may melt the insulation and may even cause fires

• This can happen in appliances with metal casing or metallic parts on the outside.

• If a kettles heating coil is touching the metal casing, the casings electrical potential increases. When you touch the kettle, there is a large potential difference between the ground and the kettle. Hence if a person touches the metal casing, a current flow through the body into the ground.

To prevent electrocution, the following precautious should be taken

• Switch off and pull out the plug from the socket when cleaning fans, television sets, computers, and toasters.

• Never insert any item into an

electrical socket. Avoid doing this even if the plastic or wooden insulator is a good insulator.

• Never overload an electrical socket.

• If electrical wiring is old and

the insulator is peeling off, the wires inside might be exposed.

• Call an electrician to repair damaged appliances or electrical cables. Do not try to fix it yourself

• Never use electrical gadgets

in wet places because water can conduct electricity through your body.

• Do not touch electrical appliances with wet hands.

• Avoid getting water into the sockets of appliances.

• Do not use an appliance until you are sure it is dry.


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Safety features in a household electrical system: Earth Wire There are three types found in the household: the live wire, the neutral wire, and the earth wire.

• Live Wire The Live Wire is (Brown/Red) is at high voltage. In

Singapore, it is at 240V. ( Most dangerous ) Fuse and switch is

located on live wire.

• Neutral Wire The Neutral Wire (Blue) is at 0V

• Earth Wire The Earth Wire (Yellow/Green) is at 0V. The earth Wire

provides a path of low resistance. The large current slows directly

from the live wire into the ground or earth. ( Must be pushed in to

open up the hold for Live + Neutral Wire )

Advantages and disadvantage of using a 2-pin plug over a 3-pin plugs?

Advantage Disadvantage 2-pin plug uses lesser material than 3-pin plug to manufacture so it economically more efficient. 2-pin plug is used for electrical appliance that do not have metal casings and hence do not need to be earthed.

Some countries have wall sockets only for 2-pin plugs, while some other countries have wall sockets only for 3-pin plugs, so an adapter is needed for the different appliances.

Fuse The fuse makes use of the heating effect of an electrical switch off a large current. A large current causes a short, thin wire in the fuse to heat up and melt. Function: The fuse is a safety device inserted in an electrical circuit to protect an equipment and wiring from excessive current flow. The fuse is always connected to the Live wire so that when a current exceeding the fuse rating flow through it, the fuse will melt and break the circuit, hence protecting the electrical appliance.

Earth Connection Symbol


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Fuse in Singapore is 13 Ampere. When the current is more than 13 Ampere, the fuse would melt causing the circuit to break. When the fuse melts, we say the fuse has ‘blown’. A blown fuse should be replaced after and electrical fault has been rectified. Circuit breaker Instead of a fuse, a circuit breaker can also cut off electric current. Unlike a fuse, the circuit breaker does not need to be replaced. When the current in one part of the circuit is too large, usually only the circuit breaker for that part turns off (Trips). Usually used in houses and buildings.


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Light There are two types of light Luminous and Non- Luminous:

Luminous objects Non-Luminous Objects • Emit Light • E.g. Sun torch, light stick,

fireflies, some types of fungi

• Reflect light • E.g. moon, table, water bottle

Luminescence = emission of light Light is an electromagnetic (EM) wave. Visible light is the range of EM radiation that is visible to human eyes. Light have both electric and magnetic field components. Properties of light:

• Light travels in straight lines – Rectilinear propagation of light.

Evidences: Ray of light in a dusty room, Solar eclipse, Lunar eclipse,

Formation of shadows.

• Speed of light, c = 3.0 X 108 m/s

• Light can travel without a medium i.e. through a vacuum

• Light consists of 7 colours – ROYGBIV

• Light is an electromagnetic (EM) wave

Electromagnetic spectrum

Gamma Rays are used for cuing cancer through radiotherapy and chemotherapy. (Extra info) Shadows

• A shadow forms when light is blocked by an opaque object.

• A solar eclipse is a phenomenon which involves the formation of

shadows.

• This happens when the moon passes between the sun and the earth.

• A solar eclipse usually occurs twice in a year.


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Reflection When light falls on a surface, several outcomes can happen:

• Only when light is reflected off the surface into our eyes can we

see the object.

• When light is completely absorbed by or transmitted (i.e. no light

reflected), the object cannot be seen.

• Light falls on an opaque object, the light energy can be absorbed

by the object and reflected.

• Light falls on a transparent object, the light energy can be

absorbed, reflected and transmitted through the object.

Regular reflection – Regular reflection occurs on smooth plane surfaces

Irregular reflection - Irregular reflection occurs on irregular or rough surfaces

Properties of a plane mirror image:

1. Virtual

2. Upright

3. Same size as object

4. Laterally inverted (Right hand side is the left hand side in the mirror)

5. Same distance behind mirror as object is in front of the mirror

2nd Law: The incident ray, reflected ray and normal at the point of incidence all lie on the same plane (coplanar). A coplanar is a plane is referring to a flat 2D surface where all these items lie on it.

1st Law: The angles of incidence and reflection are equal

i.e. i = r


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Concave Convex

• Image is enlarged and has a narrower field of vision

• E.g. Cosmetic mirrors, Dentist mirrors and Earscope

• Image is seen to be smaller than object and has a wider field of vision

• E.g. Car mirrors, Traffic mirrors and Security Mirrors

Refraction Refraction is the bending of light when light travels from a medium to another of different optical density, due to the change in the speed of light

• When light travels from less dense to denser media (slower), light

slows down and bends towards the normal. Hence the angle of

refraction is smaller than angle of incidence.

• When light travels from denser to less dense media (faster), light

speeds up and bends away from the normal. Hence the angle of

refraction is bigger than angle of incidence.

F.A.S.T. Faster away from normal Slower towards normal


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The constant value obtained shows that when light travels from air

through a specific transparent medium (say water, or glass), it always

slows down at the same rate. This constant is also known as the

Refractive Index.

• Refractive index (n) of air is 1.0003, it means light is slowed down by

1.0003 times to 2.9991x108 m/s.

• n of glass is 1.5, means light is slowed by 1.5 times to 2x108 m/s.

Speed of light changes in different mediums.

• Air - 1.00

• Glass - 1.33 Denser

• Water - 1.50

• Refractive Index, n, is a constant value unique to the transparent

medium.

• n tells us how much light is slowed down when light passes through

that medium.

• Where c is speed of light in vacuum and v is speed of light in the

medium

r

in air

sin

sin=

1st Law: The incident ray, refracted ray and normal, at the point of incidence, all lie on the same plane (coplanar).

2nd Law: Snell’s Law – For two given

media, the ratio of is a constant, where i is the angle of incidence, and r is the angle of refraction

sin i

sin r


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Dispersion of light • Dispersion is the separation of white light into its component

colours (ROYGBIV) when light travels through a lens or a prism.

• Shorter wavelengths (like violet or blue light) bend the most. Longer

wavelengths (like orange or red light) bend the least.

• All colours of light travel at the same speed in vacuum.

• Red light travels faster than Blue/Violet in most media e.g. through

glass. Hence, Red bends less than Blue.

Colours (Extra)

• We can see the colour of an object when the light falling on it

contains the same colour as the object. This is because objects

absorb some colours of light while reflecting others.

Under white light: A red, a blue and a green ball will reflect their respective colours.

The colour of light that is reflected into our eyes determines the colour of the object that appears to us.


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Measurement

• Due to the fact that there can never be a perfect measurement, data must be recorded to the correct precision according to the type of recording instrument

• There are two ways to evaluate how close the data is to the real result, and both must be present for a close tabulation:

o Precision refers to how exact a single measurement is taken (E.g. 0.1cm accuracy of ruler is less precise than 0.01cm accuracy of Vernier caliper) ▪ May be inadequate due to limitations of the instrument

o Accuracy refers to how close the reading is to the “true” value measured (E.g. Reading of 2.35 is closer to actual value of 2.30 compared to 2.00) ▪ May be inadequate due to human error or technical

skills like parallax error – minimised by repetition of experiment


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Errors

• Systematic errors o Due to error in measuring instrument (E.g. zero error) o Due to the instrument being read wrongly consistently by

person conducting experiment • Random errors

o Due to unpredictable environmental or human factors Recording of Data

• The precision of data recorded should depend on what type of instrument was used – it literally depends on the size between each measurement mark

o For example, instruments with larger spaces between each mark like ammeter, voltmeter, thermometer and measuring cylinder are recorded to half of the smallest division

o With instruments with smaller such spaces, such as rulers, protractors, electronic balances and stopwatches, precision is taken to the smallest division ▪ In particular, in stopwatches, 0.3s must be deducted to

compensate for human error Tabulation of Data Rules:

• The headings of tables should have units, so that the data does not require units

• The heading and the unit should be segregated by a / and not ( ) o E.g. L / cm, not L (cm)

• The independent variable column should increase at regular intervals

o E.g. 5, 10, 15, 20 not 6,9, 14, 21 • The precision of data should reflect what instrument was used

(follow the rules above) Recording of Calculated Data Rules:

• + or - → to smallest d.p. in equation • × or ÷ → to smallest s.f. in equation


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Graphs Remember the following:

• Scale • Title (Graph of (Y/unit) against (X/unit) • Label axes with symbol and units • Plot points with cross • Draw best-fit line

year 2 psci notes 2018 · molecules (simple and giant molecular structures) [giant is just fyi]...

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