Chemistry

Sections 5.1, 5.2, 5.5, 5.6, 5.7, 5.8, 5.9, 5.11, 5.12, 5.13, 6.5, 6.6, 6.7,

6.19, 7.1, 7.2, 7.3, 7.5,

An introduction to chemistry

• Chemistry can be defined as the study of chemicals and their reactions.

• Chemicals may be described by their physical characteristics or their chemical characteristics;– Physical characteristics include things like colour,

state at room temp., smell, boiling or melting points.– Chemical characteristics mean how a chemical

reacts with other chemicals. A chemical change occurs when a substance changes to a new substance.

Mixtures

• Most chemicals exist in nature as mixtures, made up of 2 or more substances.

• These mixtures may be either homogeneous or heterogeneous.– Homogeneous mixtures are those in which the

components are not distinguishable, is completely uniform. Ex coffee or chocolate ice cream

– Heterogeneous mixtures are those in which the components are distinguishable. Ex rocky road ice cream, stew

• Homogeneous mixtures- • Heterogeneous mixtures-

Pure Substances

• Are not as common as mixtures, consist of elements or compounds

• Elements are the simplest form of matter that can exist under natural conditions. Ex. Hydrogen, carbon, sodium

• Compounds are pure substances that contain two or more different elements in fixed proportions.

• Compounds are usually identified with a chemical formula, a combination of letters and numbers to tell you what type and how many of each element is present. Ex. H2O

H2O

WHMIS

• Stands for Workplace Hazardous Materials Information System

• Is a system to inform those using or exposed to chemicals the hazards they may encounter.

• Every chemical used in the school (cleaners included!) comes with a MSDS (Material Safety Data Sheet) that describes hazards associated with the chemical, disposal procedures etc.

• Complete questions 1, 4, 10 – 12 on page 175

ELEMENTS AND THE PERIODIC TABLE

Section 5.5

The periodic table• Organizes elements according to their atomic

structure, physical and chemical properties. • The columns (up and down) are known as

groups and the rows (across) are known as periods

• Chemical families are groups of elements that have similar properties

• We can use the organization of the elements in the periodic table to predict their reactivity (how well an element will react)

The Periodic Table

• Interactive Periodic Table

• Periodic Table: Groups and Trends

ElementsAn element is a substance made up of only 1 type of atom. There are about 112 different elements that make up the periodic table of the elements.

On the periodic table each atom type has its information. For example…

Atomic no.

Symbol

Name

Mass no.

Periodic TableAny atom can be identified by the atomic no., the symbol or by the name. For instance...

Atomic no.

Symbol

Name

Name Symbol Atomic No.

Hydrogen

Fe

12

Iron

Magnesium

H

Mg

1

26

The information from the table can also be shown as:11

5B

Questions pg 184- 186

1. Using table 1 on page 185, compare metals to non metals.

2. Where can metals and nonmetals be found on the periodic table.

3. Describe the four chemical families of the periodic table.

4. Fill in the following table about sub atomic particlesParticle Location Charge Symbol

Proton

Electron

Neutron

What it meansThe Atomic Number:

= number of protons

= number of electrons (as an atom has the same of each)

The Mass Number:

= number of protons + neutrons - why are electrons not included in the mass no?

So for Boron…

Protons =

Electrons =

Neutrons =

5

5

5.811

What about Phosphorus?

Protons =

Electrons =

Neutrons =

15

15

16

11

5B

Electron configuration

• Electrons travel in orbits or orbitals around the nucleus. The atomic number on the periodic table tells you how many electrons each element has.

• Because atoms are electrically neutral, the number of electrons equals the number of protons.

ELECTRON ARRANGEMENTElectrons are very fast moving. They are arranged in shells around the nucleus. The first shell fits…

The second fits…

The third fits…

So the electron shell for 12Mg would be…

2 e

8 e

8 e

Interactive periodic table

2, 8, 2

Ionic Bonding

• Na + Cl

2,8,1 2,8,7

IONIC FORMULAESo Mg2+ will be attracted to Cl-.

Because Mg is 2+ and Cl is only 1-, Mg will attract 2 Cl’s.

The compound formed will be MgCl2. The subscript shows that there are 2 Cl’s for each Mg.

If the starting ions were Cu2+ and S2-, the 2 ions have the same charge. So each Cu will only attract 1 S.

The compound formed will be CuS.There are never any charges on the final product - they balance out

• The mass number tells you the mass of the element and when rounded to the nearest whole number can be used to determine the number of neutrons inside the nucleus.

• Mass number – atomic number = # of neutrons.

• Ex. Oxygen Atomic # = 8Mass # = 16

8 p +8 n º

2 e6 e

Ions

• Elements are most stable where their outer electron shell or orbit is full.

• Elements whose orbit are almost full lose or gain electrons, and become ions to achieve stability

• Elements that gain electrons (and therefore a negative charge) form anions.

• Elements that lose electrons (and therefore have a positive charge) form cations.

Anions

• Are formed when non-metals gain electrons.• What was once a neutral atom becomes a

negatively charged ion.• The value of the charge is equal to the number

of electrons gained.

Cations

• Are formed when metals lose electrons • What was once a neutral atom becomes a

positively charged ion.• The value of the charge is equal to the number

of electrons lost

Naming Ions

• Cations are named by simply stating the element from which it forms followed by the word “ion”– Ex. Sodium ion

• Anions are named by stating the elements from which it forms and replacing the ending with “ide”– Ex. Chloride

Compounds: Ionic Bonding

How do atoms become stable ions?Ionic bonding animation

Types of Ions…

• Anions…Number of electrons is greater than the number of protons

• Negative charge

• Cations…number of electrons is less than the number of protons.

• Positive charge

Determining ion…general guidelines

• Metals form cations

• Non-metals form anions

Writing formulas…

• Five step rule…• 1. Write the symbol.• 2.Write the charges.• 3. Cross over the charges from top to bottom.• 4. Remove the charge.• 5. Simplify the numbers.• Formulas...

Predicting Ionic ChargesGroup 1: Lose 1 electron to form 1+ ions

H+ Li+ Na+ K+

Predicting Ionic ChargesGroup 2: Loses 2 electrons to form 2+ ions

Be2+ Mg2+ Ca2+ Sr2+ Ba2+

Predicting Ionic Charges

B3+ Al3+ Ga3+Group 3: Loses 3 electrons to form 3+ ions

Predicting Ionic ChargesNeither! Group 13 elements rarely form ions.

Group 4: Lose or gain 4 electrons?

Predicting Ionic ChargesN3-

P3-

As3-

Nitride

Phosphide

Arsenide

Group 5: Gain 3 electrons to form 3- ions

Predicting Ionic ChargesO2-

S2-

Se2-

Oxide

Sulfide

Selenide

Group 6: gain 2 electrons to Form 2- ions

Predicting Ionic ChargesF1-

Cl1-

Br1-Fluoride

Chloride

Bromide

I1- Iodide

Group 7: gain 1 electron to form 1- ion

Work for today...Chapter 5.5

• 1. Describe the alkali metals. • 2. How are the alkali metals different from the

alkali earth metals?• 3. Describe the noble gases.• 4. Describe the halogens.

Chapter 5.6Do #’s 1, 2, 3, 4Chapter 5.8Do #’s 1,2,3,4,5,6

Predicting Ionic Charges

Iron(II) = Fe2+ Iron(III) = Fe3+

Many Transition metals have more than one possible ionic charge

Predicting Ionic Charges

Zinc = Zn2+ Silver = Ag+

Some transition elements have only one possible charge

Writing Ionic Compound FormulasExample: Barium nitrate

1. Write the formulas for the cation and anion, including CHARGES!

Ba2+ NO3-2. Check to see if charges are balanced.

3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

Not balanced!

( ) 2

Writing Ionic Compound FormulasExample: Ammonium sulfate

1. Write the formulas for the cation and anion, including CHARGES!

NH4+ SO4

2-2. Check to see if charges are balanced.

3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

Not balanced!

( )2

Writing Ionic Compound FormulasExample: Iron(III) chloride

1. Write the formulas for the cation and anion, including CHARGES!

Fe3+ Cl-2. Check to see if charges are balanced.

3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

Not balanced!

3

Writing Ionic Compound FormulasExample: Aluminum sulfide

1. Write the formulas for the cation and anion, including CHARGES!

Al3+ S2-2. Check to see if charges are balanced.

3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

Not balanced!

2 3

Writing Ionic Compound FormulasExample: Magnesium carbonate

1. Write the formulas for the cation and anion, including CHARGES!

Mg2+ CO32-2. Check to see if charges are balanced.

They are balanced!

Writing Ionic Compound FormulasExample: Zinc hydroxide

1. Write the formulas for the cation and anion, including CHARGES!

Zn2+ OH-2. Check to see if charges are balanced.

3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

Not balanced!

( )2

Writing Ionic Compound FormulasExample: Aluminum phosphate

1. Write the formulas for the cation and anion, including CHARGES!

Al3+ PO43-2. Check to see if charges are balanced.

They ARE balanced!

Naming Ionic Compounds(continued)

Naming Ionic Compounds(continued)

• - some metals form more than one cation• - use Roman numeral in name

• PbCl2

• Pb2+ is cation

• PbCl2 = lead(II) chloride

• Roman numeral is equal to the charge of the cation

Metals with multiple oxidation states

Complex Ions (Polyatomic Ions)

• Mg 2+, I-, Li +, S2-

– are all called simple ions or monatomic ions

• Complex, or polyatomic ions, are tightly bound groups of ions that behave as a unit and carry a charge. Example : sulfate ion. A sulfate ion is composed of 1 sulfur atom and 4 oxygen atoms. These 5 atoms together form a unit with a charge.– SO4

2-

• Recognizing complex ions is a key in naming chemical compounds and writing chemical formulas. Polyatomic Ion Rap...

Polyatomic Ions

• Ammonium……………...• Nitrate……………………• Permanganate…………. . • Chlorate…………………• Hydroxide……………….• Cyanide………………….• Sulfate…………………...• Carbonate……………….• Chromate………………..• Acetate…………………..• Phosphate……………….

NH4+

NO3-

MnO4-

ClO3-

OH-

CN-

SO4 2 -

CO32-

CrO42-

C2H3O2-

PO43-

• cobalt (III) carbonate – Co2(CO3)3

• beryllium nitrate – Be(NO3)2

– Polyatomic tutorial...

– Please do #’s 1,2,3,4,6,7 on page 189 Chapter 5.9

Section 5.11

Molecular Compounds

Properties of Molecular Compounds

• Composed of 2 or more non-metals • Form covalent bonds in which the electrons

are shared (not lost or gained… friendlier!)• Many are gases at room temperature, they do

not conduct electricity and most are not soluble in water.

• Ionic and covalent bonds

• Uses Prefixes– 1 — mono 2 — di 3 — tri – 4 — tetra 5 — penta 6 — hexa – 7 — hepta 8 — octa 9 — nona – 10 — deca

• Example: CCl4 — carbon tetrachloride – Try: 1. P2O5

2. N2O

3. ICl3

1. Diphosphorous pentaoxide 2. Dinitrogen monoxide 3. Iodine trichloride

Naming Molecular Compounds

Name These

• N2O

• NO2

• Cl2O7

• CBr4

• CO2

• BaCl2

• H2O

• Dinitrogen monoxide• Nitrogen dioxide• Dichlorine heptoxide• Carbon tetrabromide• Carbon dioxide• Barium chloride• Dihydrogen monoxide

Write Formulas for These

• Diphosphorous pentoxide• Tetraiodine monoxide• Sulfur hexaflouride• Nitrogen trioxide• Carbon tetrahydride• Phosphorous trifluoride• Aluminum chloride

• P2O5

• I4O

• SF6

• NO3

• CH4

• PF3

• AlCl3

• To illustrate how the bonding occurs between two non-metals, Lewis Dot Structure is used.

• Only the valence electrons are used and each valence electron is represented by a dot

• Write the symbol for the element, then draw dots around the symbol to represent the number of valence electrons.

• Electrons are placed one on each side going around the symbol.

Covalent Bonding

It is only the electrons in the outermost orbits that can form bonds.

Diatomic molecules are molecules that have only 2 atoms of the same element.

• They prefer to share electrons in covalent bonds than to exist on their own.

• Element Chemical Symbol Formula and State• hydrogen H H 2 (g)

• oxygen O O 2 (g) (g) = gas

• nitrogen N N 2 (g) (l) = liquid

• fluorine F F 2 (g) (s) = solid

• chlorine Cl Cl 2 (g)

• bromine Br Br 2 (g)

• iodine I I 2 (g)• “I have no bright or clever friends”

• There is no need to balance the charges (there are no ions remember!)

• Simply look at the prefix used in the name to determine how many atoms of each element is present.

• Ex. Carbon tetrachloride C – 1, Cl – 4 CCl4

• Ex. Pentaphosphorous Trisulfide , P – 5, S – 3 P5S3

Writing Chemical Formulas for Molecular Compounds

• When using the prefixes to illustrate the number of atoms in the compound, the rules are clear;– NEVER use the prefix mono on the first element– All compounds , whether ionic or molecular, end

in “ide”– If by adding the prefix you create a double vowel,

drop the first for ease of pronunciation.

– Please do the following questions on page 204…1,2,3,4,5,6

Acids and Bases…

Acids…• Acids are sour-tasting, water soluble substances found

in many common products.• They are very reactive and good conductors of

electricity. • All acids contain hydrogen atoms in combined form and

when dissolved in water they release H+.Examples of common acids… Vinegar (acetic acid) Salad dressing Citric acid oranges, lemons Acetylsalicylic acid (ASA) Aspirin Sulfuric acid car batteries Carbonic acid carbonated drinks

Acids

• A dilute acid has lots of water and a small amount of acid

• A concentrated acid has lots of acid and not much water so must be handled carefully

• A strong acid releases lots of H+

• A weak acid releases fewer H+

Bases…• Bases are bitter tasting, water soluble and feel

slippery.They release hydroxide ions (OH-) when dissolved in water and are good conductors of electricity.

• Examples of bases…• Sodium hydroxide drain cleaner• Potassium hydroxide soap, cosmetics• Aluminum hydroxide antacids• Sodium bicarbonate baking soda

• In our home we often use bases to clean things… Bleach and toothpaste

• Some things are not acids or bases: we say that they are neutral…eg. water

Recognizing acids and bases from their chemical formulas…

• 1. Acids are easy! They begin with hydrogen H2SO4 – sulfuric acid, or H2CO3 – carbonic acid.

• 2. Bases are more difficult. They usually contain OH but not always, ex. NaOH. An exception would be NaHCO3 (baking soda) is a base because it reacts with water to produce Oh-.

• Questions…1. What is the most important acid in the chemical industry?2. What is it used for?3. #3 on page 295

Chapter 6 - Understanding Chemical Reactions

• A word equation is one way of representing a chemical reaction. It tells you what reacts and what is produced.

• Word equations are written like this:reactants products

DO NOT COPY: When hot steel wool (iron) is put into a bottle of oxygen,

there is a spectacular reaction and iron (III) oxide is produced. The word equation would be:

iron + oxygen iron (III) oxide

• Write the word equation for the following example:

When zinc is added to hydrochloric acid, hydrogen and zinc chloride are produced.

Zinc + hydrochloric acid hydrogen + zinc chloride

• Please do questions 2 and 3 on page 219

6.5 Balancing Chemical Equations

• A skeleton equation represents all chemicals by their formulas.

Word Equationmethane + oxygen water + carbon

dioxideSkeleton Equation

CH4 + O2 H2O + CO2

6.3 Conserving Mass

• The Law of Conservation of Mass states that in a chemical reaction, the total mass of the reactants is always equal to the total mass of the products.

How to balance Equations:• 1. Determine the correct formulas and write

the skeleton equation:Fe + O2 Fe2O3

• 2. Count the number of atoms of each element in the reactants and products. (Polyatomic ions appearing unchanged on each side are counted as a single unit).

Type of Atom Reactants ProductsFe 1 2O 2 3

• 3. Balance the elements one at a time by using coefficients. The coefficient is a whole number that appears in front of the formula. When no coefficient is written, it is assumed to be 1.

• 4 Fe + 3 O2 2 Fe2O3

• Always check to be sure that the equation is balanced.

• Type of Atom ReactantsProducts

• Fe 4 4• O 6 6

5. Make sure all coefficients are in the lowest possible ratio. (Reduce if possible)

Balancing Chemical Equations

Section 6.5

• Because we cannot change the chemical formulas of compounds in the reaction, we need to use coefficients to balance the number of atoms.

• Coefficients are numbers placed in front of the compound and apply to all elements in the compound (unlike subscripts which only apply to that element).

Equation types

• We began by writing word equations– Iron + Oxygen Iron II oxide

• Writing chemical formulas based on the word equations is known as skeleton equations.– Fe + O2 FeO

• Because of the Law of Conservation of Mass, we now need to write a balanced equation.

Steps to balancing equations

1. Count the number of atoms of each type in the reactants and products.

Fe + O2 FeO

Type of Atom Reactants Products

Fe 1 1

O 2 1

2. Multiply each of the formulas by the appropriate coefficient to balance the number of atoms. Re-write the equation.2Fe + O2 2 FeO

Tips• Look for larger molecules (ie polyatomic) or

complex molecules and balance them first, especially if they appear on both sides.

• If you have an odd number on one side and an even number on the other, fix the odd side first.– Al + O2 Al2 O3 , fix products first

• Leave diatomic molecules and elements that appear more than once on the same side to the end.

6.7 Types of Reactions

• There are five main categories of chemical reactions:

1. Combustion2. Synthesis3. Decomposition4. Single Displacement5. Double Displacement

• Combustion –the rapid reaction of a substance with O2 to produce compounds called oxides (often call this process burning).

• The fuel can be a variety of things but it is often a hydrocarbon (ex. gasoline) The formula for combustion of a hydrocarbon isC4 H10 + O2 CO2 + H2 O + energy

(C4H10 is Butane)

The products of a combustion reaction are always carbon dioxide and water. (C4H10 is butane)

Skeleton...C4 H10 + O2 CO2 + H2 O + energy

Balanced ...

2C4 H10 +13O2 8CO2 + 10H2 O + energy

Synthesis Reactions

• Involves the combination of smaller atoms or compounds into larger compounds. (also known as combination reactions).

• They have the following general formula:

• A + B AB

• If both reactants are elements then the reaction MUST be synthesis.

• Example: 2H2 + O2 2H2O

Examples

• 2Na + Cl2 => 2NaCl

• 2Al + 3Br2 => 2AlBr3

• Synthesis reactions sometimes involve joining two compounds into a larger one.

hydrogen chloride + ammonia ammonium chloride

HCl + NH3 NH4Cl

Decomposition Reactions

• Involves the splitting of a large compound into smaller molecules or elements.

• They have the following general formula: • AB A + B

• If there is only 1 reactant then the reaction MUST be decomposition.

• Example: 2H2O 2H2 + O2

What types of Reactions are these?

1. H2CO3 CO2+ H2O

2. 2Fe + O2 2FeO

3. C10H8 + 12 O2 10 CO2 + 4 H2O

Answers:1. Decomposition2. Synthesis3. Combustion

Please do the following…

• Page 235…#’s 1,2,3,4

Single Displacement

• This is when one element trades places with another element in a compound. These reactions come in the general form of:

A + BC ---> AC + B• Example: Fe + CuSO4 => FeSO4 + Cu

• The reactants MUST be an element and a compound

• Single displacement can involve metals:Na + KCl K + NaCl

• Single displacement can involve nonmetals:F2 + 2LiCl 2 LiF + Cl2

• Remember - If the single element is a nonmetal it will replace the nonmetal.

• If the single element is a metal it will replace the metal.

Double Displacement

• Involves two elements replacing one another.

• The reactants must be compounds (usually happens in solution).

• The positive ions stay in the same position (A and C) and the negative ions change partners (B and D). The general formula is:

AB + CD AD + CB

• NaOH + FeCl3

NaOH + FeCl3 Fe(OH)3 + NaCl

• Pb(NO3)2 + 2 KI

Pb(NO3)2 + 2 KI PbI2 + 2 KNO3

List what type the following reactions are:

• 1) NaOH + KNO3 --> NaNO3 + KOH

• 2) CH4 + 2 O2 --> CO2 + 2 H2O

• 3) 2 Fe + 6 NaBr --> 2 FeBr3 + 6 Na

• 4) CaSO4 + Mg(OH)2 --> Ca(OH)2 + MgSO4

• 5) Pb + O2 --> PbO2

• 6) Na2CO3 --> Na2O + CO2

• 1) double displacement 2) combustion 3) single displacement 4) double displacement 5) synthesis

6) decomposition

Please do the following…

• Page 241…#’s 1,2,3

• Chemical reactions and balancing equations...

Rates of reaction

Objectives• To understand that a chemical reaction

involves collisions between particles• To be able to describe the four factors which

will affect the rate of a chemical reaction.

How do we make the reaction go faster?

• There are four things that we can change to make the reaction go faster.

They are: • Temperature• Surface area• Concentration• Using a catalyst

Temperature

• When we increase the temperature we give the particles energy

• This makes them move faster • This means they collide with other

particles more often• So the reaction goes faster.

Surface area

• If we make the pieces of the reactants smaller we increase the number of particles on the surface which can react.

• This makes the reaction faster.

The particles on the surface can react

When cut into smaller pieces the particles on the inside can react

Concentration• If we make one reactant

more concentrated (like making a drink of orange squash more concentrated)

• There are more particles in the same volume to react

• So the reaction goes faster.

There are less red particles in the same volume so there is less chance of a collision

There are more red particles in the same volume so there is more chance of a collision so the reaction goes faster

Using a catalyst

• A catalyst is a chemical which is added to a reaction.

• It makes the reaction go faster.• The catalyst does not get used up in the

reaction.• It gives the reaction the energy to get

started

Click here ...to complete exercise 2

Click here ...to complete exercise 3

Rates of reaction