CHEMISTRY noun \ˈke-mə-strē\ - the study of matter & its changes

Matter & Energy comprise every known thing in the universe.

• Matter - anything that has mass and takes up space

• Mass – the amount of matter in something

Measured in grams (g) with a triple-beam balance

• Volume – the amount of space something occupies

Measured in milliliters (mL) with a graduated cylinder

Matter

Atoms

The smallest unit of matter is an atom, which are so small they cannot be seen even with a microscope.

Element

An element is a pure substance that cannot be broken down into any other substance by chemical or physical means. Examples:

•aluminum

•zinc

•oxygen

Subatomic Particles

Each element has it’s own unique atom that has a unique number of subatomic particles.

In the nucleus (center):Protons – positive chargeNeutrons – are neutralOrbiting the nucleus :Electrons – negative

charge

The Periodic Table

The periodic table contains information about all the different elements that make up our universe.

Atomic Number

Elements are organized on the periodic table by their ATOMIC NUMBER, which is the number of PROTONS in the nucleus.

Isotopes Atoms of the same element

can have different numbers of NEUTRONS; the different possible versions of which are called ISOTOPES.

Neutrons

To determine the average number of neutrons in an element: Round the atomic mass to the nearest whole number and subtract the atomic number (of protons).

Ex: Li 7 – 3 = 4 neutrons

Atomic Mass

The ATOMIC MASS listed on

the periodic table is an average of the mass of all known isotopes of that element.

Electrons

Electrons orbit the nucleus in an electrons cloud.

There are up to 7 ENERGY LEVELS within the electron cloud and each can only hold a certain number of electrons.

Valence Electrons

The electrons on the outermost shell of an atom that participate in bonding are called valence electrons.

Symbols

Hg – Mercury Au – Gold Pb – Lead Sn – Tin Ag – Silver Cu – Copper Fe – Iron K – Potassium Na - Sodium

Patterns

The PROPERTIES of an element can be predicted by its location on the periodic table.

Periods & Groups

The period (row) on which an element can be found will tell you how many energy levels are needed to house all electrons.

Groups / Families

There are 18 groups from left to right.

Like family members,

elements of the same group have similar chemical properties.

Group 1

ALKALI METALS• Hydrogen is not a member, it is a non-

metal• All are metals and solid at room temp• 1 Valence Electron• Soft and silvery, shiny• Very reactive, esp. with water• Conduct electricity

Group 2

Metals

Solids at room temp

• 2 electrons in the outer shell

• White, silvery, and malleable

• Reactive, but less than Alkali metals

• Conduct electricity

Groups 3-12

TRANSITION METALS• Metals that provide many colorful

pigments & durable building materials.• Almost all are solids at room temp (Hg)• Good conductors of heat & electricity.• 1 or 2 Valence Electrons• Less Reactive than Alkali and Alkaline

Earth• Can bond with many elements in a

variety of shapes.

The Lanthanide & Actinide Series of Periods 6 & 7

• Some (mostly man-made)are Radioactive

• The rare earths are silver, silvery-white, or gray metals.

• Super conductors - Conduct electricity very well

Rare Earth Metals

Metalloids

Along the stair step line – (7) METALLOIDS that have

properties of metals and non-metals are useful because of their unique properties.

Ex: silicon semi-conductor for computer chips

Non-Metals

• The NON-METALS are not good conductors but they combine with others readily to form compounds.

• Many of these elements are crucial in creating and maintaining life (C, N, O, P, S).

• They are found to the right of the stair-step line.

Group 17

7 electrons in the outer shell• Non-metals, Uus is unknown• Very reactive - are often bonded with

Group 1 Alkali Metals• Has 2 gases, 1 liquid (Br), and 2 solids• Diatomic – meaning they are never

found as a solitary atom.

Group 18

• Exist as gases• Non-metals• 8 electrons in the outer shell = Full• Helium (He) has only 2 electrons in

the outer shell = Full• Not reactive with other elements

Natural vs. Synthetic

The majority of the known elements are naturally occurring, however all elements above 92, are known as: SYNTHETIC elements having been created in a lab.

Chemical Bonding

Knowing the number of valence electrons allows us to make predictions about how they will combine / bond with other elements to make molecules.

Molecules 2 or more atoms help

together by chemical bonds

Covalent Bonds

A chemical bond that involves the SHARING of their valence electrons.

non-metal + non-metal

Ionic Bonds

A bond that involves the TRANSFER of valence electrons from one atom to another.

metal + non-metal

Ions

Atoms that gain or lose electrons become heavy on the + or - charge

Cations +

Atoms with a positive

charge. These are metals since they lose electrons because they have one or 2 valence electrons to spare.

Anions -

Atoms with a negative

charge. These are nonmetals since they gain electrons because they only need 1 or 2 to fulfill their life’s goal of having 8 valence electrons.

Compound

A compound is a pure substance made of 2 or more elements chemically combined in a set ratio. Compounds cannot be easily separated.

Mixture

A mixture is a pure substance made of 2 or more elements, compounds, or both, that are together yet not chemically bonded and therefore can be separated by physical means.

Heterogeneous Homogeneous

Solutions

Solutions are homogeneous mixtures in which one substance (the solute) is dissolved in another (the solvent).

It has the same properties throughout and contains

solute particles (molecules & ions) that are too small to see.

Solutions with Water

Water is often referred to as the “universal solvent” because there are a wide variety of things that will dissolve in it.

Ex: sugar + CO2 + H2O makes sodabodily fluids (for plants and animals)

Types of Solutions

A solution can be a:•Solid dissolved in a liquid - salt water•Liquid dissolved in a liquid - antifreeze•Solid dissolved in a solid – alloys•Gas dissolved in a liquid – soda

•Gas dissolved in a gas – air N2 & O2 +

Not Solutions!

Suspension – A mixture in which the particles can be seen and easily separated using settling or filtration.Ex: pepper in water, sand in waterColloid – A mixture that contains small, undissolved particles that do not settle out and are large enough to scatter a beam of light.Ex: fog, milk, mayonnaise, shaving cream, whipped cream

Particles in Solution

Some compounds break into their ions in solution (table salt - Na+Cl-) While others break into neutral molecules (sugar – C12H22O11)

For this reason, some solutions CONDUCT electricity (ionized) While other (neutral ones) do not.

Effects of Solutes on Solvents

Solutes can lower the freezing point and increase the boiling point of solvents.

Concentration

Concentration is the ratio of solute to solvent.

A dilute solution has a small amount of solute.

A concentrated solution has a large amount of solute.

Solubility

Solubility is the amount of solute that can be dissolved in a solvent at a given temperature.

A saturated solution is solution in which so much solute has been added that it will no longer dissolve.

An unsaturated solution will continue to dissolve solute.

Identification by Solubility

Solubility is a PHYSICAL property, therefore it can be used to help determine the identity of a mystery substance in a lab!

Example

Factors Affecting Solubility

Several factors affect solubility:•Pressure – in gases, a gas under higher pressure can dissolve more solute.•Type of Solvent – like dissolves like•Temperature – more solute can be dissolved at a higher temperature

Supersaturation

A heated solution can typically dissolve more solute than it does under normal conditions.When more solute is added under these conditions and then the solution is cooled, the solution is said to be supersaturated.Ex: good ole southern sweet teaAdding a tiny additional bit of the solute at that point will cause it all to fall out of solution.

Mixture vs. Compound

Properties

A property is a characteristic of a substance that can be observed.

Physical Properties

A physical property is one that can be observed without changing the identity of the substance.

Examples• Malleability: the ability to be hammered into

a thin sheet

• Ductility: the ability to be stretched into a wire

• Melting/freezing point: solid -> liquid <-solid

• Boiling point: liquid -> gas

• Density: mass/volume

• Solubility

• Specific heat: the amount of heat required to heat a substance 1 degree Celsius

• Luster: shiny, matt

Density

• Density is the amount of mass per unit of volume.

• Like many other properties it can be used to identify a substance.

• The density of water is 1.0g/mL

Calculating Density

D = mass =_g_ = _g_ volume mL cm3

Ex: A cube has a mass of 2.8 g and occupies a volume of 3.67 ml. Would this object float or sink in water?

Mass = 2.8 g Volume = 3.67 mL D = 2.8g/3.67 mL= 0.76 g/mL

Identification by Density

A liquid has a mass of 25.6 g and a volume of 31.6 mL.

Use the table below to identify the substance.

M=25.6 g V=31.6 mL

D = 25.6 g/31.6 mL

D= 0.81 g/mL

The substance is ethyl alcohol.

Chemical Properties

A chemical property is a property that can only be observed by changing the identity of the substance

Examples:

•flammability

•ability to rust

•reactivity with vinegar

Changes

• Physical change – substance maintains its chemical makeup

Ex: state changes, dissolving• Chemical change– substance

becomes something else entirely Ex: burning, oxidation

Chemical Change

Chemical change occurs when bonds break and new bonds are formed.

The chemical composition (makeup) of the substance(s) has been altered and it is no longer the same substance.

Chemical Reactions

• Reactants – are the chemicals that go into a reaction.

• Products – are the chemicals products that are created by the reaction.

Law of Conservation of Matter

States that the:Mass of the reactants = Mass of the Products(in a closed system)

Matter can be neither created nor destroyed!

Ex: 5 g of sodium(Na) + 5 g of chloride (Cl) yields 10g of table salt (NaCl)

Chemical Equations

Reactants Products 2H2 + O2 2H20

Coefficient – # of moleculesSubscript - # of atoms

Balancing Equations

In order to accurately demonstrate the Law of Conservation of Matter you MUST have the same number of atoms of each element on both sides of the reaction.

Watch this video!

Types of Chemical Reactions

• Synthesis A + B C• Decomposition C A + B• Replacement AB + CD AD + BC

Evidence of Chemical Change

• Change of properties • heat absorbed -endothermic• heat released – exothermic

• gas formation (O2, CO2) – “bubbling”

• Precipitate formation - a solid formed from 2 liquids

Exothermic

Produces Heat

Endothermic

Soaks up heat from the surroundings; observed as a decrease in temperature

Activation Energy

Activation Energy is the minimum amount of energy required to start a chemical reaction.It has a cascading effect.An endothermic reaction requires A LOT of activation energy.

Demonstration Reaction

Citric Acid + Sodium Bicarbonate

H3C6H5O7(aq) + 3NaHCO3(s) →

3CO2(g) + 3H2O(l) + Na3C6H5O7(aq)

Follow Up Questions

Answer on a sheet of notebook paper:1. Describe what happened in this

demonstration.2. What was the initial temperature in

the beaker?3. What was the final temperature in

the beaker?4. Was this an exothermic or an

endothermic reaction? 5. What type of reaction was this? Synthesis, decomposition or replacement

Surface Area

If you break the reactants into smaller pieces then there is more surface area in contact with the other reactant. Thereby increasing the chance that 2 oppositely charged atoms can bond together.

Temperature

If you increase the temperature then the molecules are moving faster and thereby making connections more frequently.

Concentration

Concentration is the amount of a substance in a given volume. Increasing the concentration of a substance means there are more atoms or molecules available for bonding.

Catalysts

Catalysts increase the rate of reaction by lowering the activation energy required to start the reaction.They are not reactants and they are not consumed during a reaction.

Inhibitors

Inhibitors slow reactions by interfering with the reactants ability to get to each other.

Demonstration Reaction

Decomposition of Hydrogen Peroxide

KI

2H2O2 2H2O + O2 + HEAT!

Follow Up Questions

Answer on a sheet of notebook paper:1. Describe what happened in this

demonstration.2. Was this an endothermic or

exothermic reaction?3. What was the catalyst in this

reaction? 4. What 2 products caused the

dramatic “smoke”?5. What type of reaction is this? Synthesis, decomposition or replacement

Precipitate Reactions

If the ions in 2 solutions combine to form a solid and that solid is NOT soluble with the solvent produced a precipitate will form.

Example

Fe(NO3)3(aq) + 3 NaOH(aq) Fe(OH)3(s) + 3 NaNO3(aq)

Demonstration Reaction

Cu(NO3)2(aq) + 2 NaOH(aq) Cu(OH)2(s) + 2 NaNO3(aq)

Follow Up Questions

Answer on a sheet of notebook paper:

1. Describe what happened in this demonstration.

2. What is a precipitate? 3. What type of reaction is

this? Synthesis, decomposition or

replacement

Acids

Produce H+ ions in H2O

Ex: HCl H+ + Cl-

Properties:•Tastes sour• Corrosive reaction w/metal

• Reacts w/CO32- to make CO2

• Turns blue litmus paper red

Examples of Acids

• HCl – hydrochloric acid

• CH3CO2H – acetic acid (vinegar)

• H2SO4 – sulfuric acid

• Ascorbic Acid – Vitamin C (citrus)• Fertilizers – Nitric & Phosphoric Acid• Lactic Acid

Bases

Produce OH- ions in H2O

Ex: NH3 + H2O NH4+ + OH-

•Tastes bitter• feels slippery• Turns red litmus paper blue

Examples of Bases

Strength

The strength of an acid or a base is based on how well it produces ions in water.

Strong Acids – HCl & H2SO4

Strong Bases - NaOH

Measuring Strength

pH = Potential Hydrogen

Range of values from 0 to 14 that describes the concentration of H+

ions in a substance.

Safety

Know the pH (strength) of the acid or base you are handling.

Everything from 2-12 is in the safe zone.

PrecautionsWhen working with a strong acid or base (0-2 or 11-14) be sure to wear goggles…even if it’s a dilute solution!

For spills-Pour vinegar on a base & sodium bicarbonate on an acid…because…

Acid Base Neutralization

Displacement reaction

Acid + Base (liquid) water +(solid) salt

Salt = Group 1-2 Metal + a halogen

Demonstration Reaction

NaHCO3 + NaOH Na2CO3 + H2O

Chemicals Everywhere

Chemistry all around us:•Beauty products•Cleaning products•Food products•Scents & flavors•Monitoring the environment•Containers•Protection •Explosives

Materials

We use the following materials on a regular basis, they are made of various substances that we’ve been studying lately:•Plastics•Metals•Alloys•Ceramics•Glass

Polymers

Large complex molecules made from smaller molecules joined together in a repeating pattern (chain).

They are both naturally occurring & synthetic.

Forming Polymers

Carbon Structures

Polymers are mainly composed of various configurations of Carbon (C) & Hydrogen (H)

Natural Polymers

• Cellulose – cell walls of fruits & vegetables

• Starches – pasta, bread & vegetables

• Natural fibers – hair, wool• Amino Acids Protein

DNA!

Synthetic Polymers

• Plastics MANY types (see p.731)• Fibers – carpets, nylon• Chewing gum• Teflon coating

Plastics

Benefits:• Cheap & easy to make• Lightweight• Versatile• DurableProblems:• Disposal Recycling can be cost prohibitive

Alloys

A mixture of 2 or more elements at least one of which is a metal.

Alloys are usually stronger & more durable than the metals which they are made from. They are also less likely to suffer corrosion from oxidation.

Examples

Ceramics

Hard, crystalline solids made from heating clay (water & minerals – Si, AL & O) to very high temperatures.

Properties

• Water resistant• Strong building material -

bricks• Not conductive of electricity• Can withstand much higher

temperatures than most metals

• Brittle & prone to breaking

Uses

• Containers • Storage• Cookware• Archaeologists use ceramic

sherds for relative dating!• Home construction – roofs &

floors• Replace joints – hips, knees,

teeth (dentures)

Glass

Molten sand & limestone can be shaped and cooled into waterproof vessels.

Uses

• Storage• Containers• Windows• Lenses for eyeglasses,

telescopes, microscopes• Cookware• Data transmission (optical

fiber)