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$: Unit 1 - Chemical Processes - Chapter 1 - Chemical ... · Chapter 1 - Chemical Nomenclature Watch: \An Introduction to Chemistry" Unit 1 ... write the chemical formulas for ionic$
Unit 1 - Chemical ProcessesChapter 1 - Chemical Nomenclature

Watch: “An Introduction to Chemistry”

Unit 1 - Chemical Processes 1 / 57

https://www.youtube.com/watch?v=izeuGr0lbN0

Part I - Chemicals in Action

This section will explore ionic and molecular chemical compounds. Wewill:

recognize the relationships among chemical formulas, composition,and names;

write the chemical formulas for ionic and molecular compounds, andname these compounds.


Part II - Understanding Chemical Reactions

Chemical reactions can be classified according to the types of reactantsinvolved. We will:

balance chemical equations;

recognize the types of chemical equations.


Part III - Controlling Chemical Reactions

This section will discuss the rates of chemical reactions. We will:

explain which factors affect rates of chemical reactions and why;

identify everyday examples of ways in which rates of reactions arecontrolled.


Section 1.1 - Introduction to Chemistry

What is Chemistry?

Chemistry is the study of the properties of materials and the changes thatmaterials undergo.


Why study Chemistry?

It is the central science, leading to a fundamental understanding of othersciences and technologies.

It is an extremely practical science that greatly impacts our daily living:

improvement to health care

conservation of natural resources

protection of the environment

provision of our everyday needs for food, clothing, and shelter

Using chemistry, we have discovefirebrick helpful pharmaceuticalchemicals, increased food production, and developed plastics.


Classifications of Matter

Matter is the physical material of the universe. The tremendous variety ofmatter in our world is due to the combinations of only about 100 verybasic substances called elements.

How can we classify matter?


Pure Substances

A pure substance is matter that has distinct properties and acomposition that doesn’t vary from sample to sample.

Examples:

Water

Oxygen

Table salt

All substances are either elements or compounds.

Elements are composed of only one kind of atom (ex: O, H, Fe).

Compounds are substances composed of two or more elements(water, table salt).


Properties of Matter

Physical property - is a characteristic of a substance.

Example: Baking soda is a pure substance that is white, solid at roomtemperature, and dissolves readily in water (ex. changes of state).

Physical Change - a change in the size or form a substance which doesnot change the chemical properties.

Freezing - liquid to solid

Melting - solid to liquid

Boiling - liquid to gas

Condensation - gas to liquid

Sublimation - solid to gas


Chemical property - a characteristic behaviour that occurs when asubstance changes to a new substance.

New substance is formed

Gas is produced

Colour change

Difficult to reverse

Heat or light is given off

Solid is formed


Section 1.2 - Chemical Safety Symbols

Hazardous Household Product Symbols (HHPS)

Toxic - substances that even insmall quantities may poison,cause injury or death whenswallowed, absorbed through theskin, or inhaled into the lungs.

Flammable - substances, usuallyliquids, that can readily ignite(burn in air) in a wide range oftemperature conditions.


Corrosive - substances orvapours that can deteriorate oreat away the surface of anothermaterial.

Reactive/Explosive -substances that can react withair, water, or another substanceto produce toxic vapours orexplode.

Watch: “Sodium in Water”


https://www.youtube.com/watch?v=RAFcZo8dTcU

Degree of Danger - combinedwith these three symbols, theprevious classification imagesshow the type and extent towhich a substance can beharmful.


Workplace Hazardous Materials Information System (WHMIS)

Class A: Compressed gas

Class B: Flammable andcombustible material

Class C: Oxidizing Materials


Class D: Poisonous and infectious materials

Division 1: Materials Causingimmediate and Serious ToxicEffects

Division 2: Materials CausingOther Toxic Effects

Division 3: BiohazardousInfectious Material


Class E: Corrosive material

Class F: Dangerously reactive material

NFPA - National Fire ProtectionAgency


Exit Slip

URL: http://goo.gl/LTRNYX

Recall:

Matter is the physical material of the universe.

All matter is composed of elements.

All elements are composed of very small particles called atoms.

What are atoms made up of?


Atomic structure:

Atoms contain two main regions called the nucleus and the electroncloud.

The nucleus is composed of two subatomic particles called protonsand neutrons.

The electron cloud is composed of one subatomic particle called theelectron.

(a) The Bohr Model (b) The Electron Cloud Model


The charges of the three key subatomic particles:

Protons - Positively charged. Denoted: p+

Neutrons - Neutrally charged. Denoted: n0

Electrons - Negatively charged. Denoted: e–

Elements on the Periodic Table are in electronically neutral form; i.e.they carry no charge. To be electronically neutral, an atom must have anequal number of protons and electrons.

Example: Which atom below is electronically neutral?


Question:

Label the following diagram:


Section 1.3 - The Periodic Table

Periodic Table - A structured arrangement of elements that allows us toexplain and predict physical and chemical properties.


Chemical Families or Groups - Elements in the same vertical column ofthe periodic table. They tend to have similar physical and chemicalproperties.


Symbols on the Periodic Table


Determining the number of protons, neutrons, and electronsin an atom:

For any element:

Number of protons = Atomic Number

Number of electrons = Number of Protons = Atomic Number

Number of Neutrons = Mass Number - Atomic Number

Examples - Calculate the number of protons, neutrons, and electrons inthe following atoms:

1 Lithium

2 Krypton

3 Sodium

4 Potassium


Section 1.4 - Bohr Diagrams: A Model for an Atom

Each shell or orbit can hold only a certain number of electrons.1 First orbit - 2 electrons2 Second orbit - 8 electrons3 Third orbit - 8 electrons4 Fourth orbit - 2 electrons

The valence shell of an atom is the outermost orbit. The electrons in thisorbit are called valence electrons.


Examples:

Draw Bohr diagrams for the following:

1 a helium atom

2 a carbon atom

Identify how many valence electrons are present in each. Include thenumber of protons and neutrons in the nucleus.


Practice:

Draw Bohr diagrams for the following:

1 a potassium atom

2 a nitrogen atom

Identify how many valence electrons are present in each. Include thenumber of protons and neutrons in the nucleus.


The Nobel Gases

The nobel gases behave differently than other elements on the periodictable. Draw Bohr diagrams for the following:

1 Helium

2 Neon

3 Argon

What do you notice about the number of electrons in each orbit?

The nobel gases do not easily form compounds because their arrangementsof electrons are very stable. Their orbits are full.


When elements have the opportunity, they will give, receive, or shareelectrons so that their electron arrangements match the most stableelectron arrangement of the closest nobel gas.

Example: How could lithium obtain an electron arrangement of itsclosest noble gas?


Ions

Consider lithium (3 electrons). If lithium loses the electron in its outerorbit, than it’s left with 2 electrons - the same stable electron arrangementas helium. However, it still has 3 protons.

=⇒ It has formed an ion. When atoms gain or lose electrons theybecome ions.

Its ionic charge is now 3 + (−2) = 1. This is written as Li1+ or Li+.


If an element loses electrons it becomes positively charged and is called acation.

Example: Li+

If an element gains electrons it becomes negatively charged and is calledan anion.

Example: F−


Examples:

1 How many electrons does Na+ have? What type of ion is Na+?

2 How many electrons does Cl– have? What type of ion is Cl– ?


Practice: What ionic charge would the following elements have to beclosest to a noble gas. How many elements would these ions have? Whattype of ions would they be?

1 Li, Na

2 Be, Mg

3 B, Al

4 C, Si

5 N, P

6 O, S

7 F, Cl


How Elements Form Compounds

Much of chemical activity involves the transfer of electrons from onesubstance to another.

Ionic compounds are generally combinations of metals and nonmetals(cations and anions). These compounds are held together by attractionsbetween opposite charged ions (like a magnetic attraction). This is calledan ionic bond.

Molecular compounds are generally composed of nonmetals only.Electrons are shared between atoms. This is called a covalent bond.

Examples:1 MgCl2 (magnesium chloride)

2 H2O (water)


Examples: Using Bohr diagrams, show the complete electron transferthat results in the formation of the following ionic compounds.

1 NaCl

2 LiF

3 CaF2

Practice:

1 KCl

2 MgCl2

3 BeO


Section 1.5 - Naming Ionic Compounds

Steps for naming binary ionic compounds:

1 Name the metal (cation) first.

2 Name the nonmetal (anion) second. For the nonmetal, drop theending and add “ide”.

3 Use Roman Numerals in parentheses to indicate which charge of themetal is used. This is only used for metals that make more than onecharge.


Examples:

1 Ca3N2

2 Mg3P2

3 Al2O3

4 CuCl2

5 FeBr3

6 CoN


Practice: Name the following compounds.

1 CaCl2

2 MgBr2

3 FeI3

4 CuCl2

5 Cr3N2

6 BeO


Polyatomic Ions

Polyatomic ions - groups of atoms that tend to stay together and carryan overall ionic charge.

Examples:

nitrate - NO –3

hydroxide - OH–

bicarbonate - HCO –3

chlorate - ClO –3

sulfate - SO 2–4

phosphate - PO 3–4

Note: They are all anions!


Steps for naming ionic compounds with polyatomic ions:

1 For polyatomic compounds, name the metal first followed by thename of the polyatomic ion. The endings do not change.

Examples:

1 Al(OH)3

2 NaHCO3

3 Sn(NO3)4


Practice: Name the following compounds.

1 KNO3

2 Mg3(PO4)2

3 NaOH

4 Pb(NO2)4


Section 1.6 - Writing Formulas for Ionic Compounds

Steps:1 Write the ionic charges above the symbols.2 Determine the lowest number of each element that would make the

entire compound electronically neutral.i.e. “Crisscross” the numbers, using them as subscripts. Remember

to reduce when appropriate.

Key fact: The sum of the charges on the positive ions equals the sum ofthe charges on the negative ions.

Examples:1 calcium iodide

2 sodium phosphide

3 calcium chloride


Practice: Write the formulas for the following compounds.

1 sodium bromide

2 strontium nitride

3 potassium phosphide

4 magnesium nitride

5 zinc iodide


Note: Some metals can make more than one kind of ion. Romannumerals in brackets indicate the charge of the ion (not how many are inthe formula).

Examples:

1 iron (III) oxide

2 iron (II) oxide



1 cobalt (II) chloride

2 nickel (III) oxide

3 lead (IV) selenide

4 tin (II) nitride

5 iron (III) phosphide


Note: Ionic compounds may include polyatomic ions. Reference the “Listof Polyatomic Ions” to determine the charges of the polyatomic ions.

Remember to put the entire polyatomic ion in brackets if you need morethan one!

Examples:

1 sodium carbonite

2 aluminum sulfate



1 sodium chlorate

2 magnesium phosphate

3 silver nitrate

4 barium sulfite

5 iron (II) nitrite


Section 1.7 - Molecular Compounds

Molecular compound - formed when nonmetals (anions) share electronswith other nonmetals (anions).

Covalent bond - a shared pair of electrons held between two nonmetalatoms that holds the atoms together.


Examples: Using Bohr diagrams, show the complete electron sharing thatoccurs in the following molecular compounds.

1 H2

2 Cl2

3 H2O


Practice: Using Bohr diagrams, show the complete electron sharing thatoccurs in the following molecular compounds.

1 F2

2 HF

3 CH4


Diatomic Molecules

These are molecules composed of two of only one type of element.

The most common are:

H2 O2 F2 Br2 I2 N2 Cl2

Acronym: HOFBrINCl

Why would these atoms bond with themselves to form covalentbonds?


Writing Formulas for Molecular Compounds

Steps:

1 Write the symbols of each nonmetal.

2 Write the appropriate subscripts for each nonmetal that correspondsto the prefixes used.

Note: The prefix “mono” is never used on the first nonmetal.

mono di tri tetra penta hexa hepta octa nona deca

1 2 3 4 5 6 7 8 9 10


Examples: Write the formulas for the following molecular compounds.

1 carbon dioxide

2 dinitrogen trioxide

3 carbon tetrafluoride

Practice: Write the formulas for the following molecular compounds.

1 dihydrogen monoxide

2 diboron trioxide

3 phosphorus trihydride


Writing Names for Molecular Compounds

Steps:

1 Write the name of the first nonmental and then the second.

2 Place the appropriate prefixes in front of each of the names. Theprefixes correspond to the subscripts of each element.

3 Drop the ending of the last element named and add -ide.

Note: Sometimes prefixes are shortened when the ending vowel of theprefix “conflicts” with a starting vowel in the compound. This makes thename easier to pronounce.


Examples: Name the following molecular compounds.

1 CS2

2 N2O4

3 P4S10

Practice: Name the following molecular compounds.

1 CO

2 CF4

3 C2H6


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