1. Demonstrate knowledge of the three subatomic particles, their properties, and their location within the atom. 2. Define and give examples of ionic bonding (e.g., metal and non‐metal) and covalent bonding (e.g., two non‐metals, diatomic elements). 3. With reference to elements 1 to 20 on the periodic table, draw and interpret Bohr models, including protons, neutrons, and electrons, of: • atoms (neutral) • ions (charged) • molecules ‐ covalent bonding (e.g., O2, CH4) • ionic compounds (e.g., CaCl2)

4. Identify valence electrons using the periodic table. 5. Distinguish between paired and unpaired electrons for a single atom. 6. Draw and interpret Lewis diagrams showing single bonds for simple ionic compounds and covalent molecules (e.g., NaCl, MgO, BaBr2, H2O, CH4, NH3). 7. Distinguish between lone pairs and bonding pairs of electrons in molecules.

Alkali earth metals

Alkali metals

Anions

Atomic #

Atomic number

Atomic Theory

Atoms

Bohr diagram

Cations

Chemical Change

Chemical reaction

Compound

Covalent bonding

Molecule

Neutron

Noble gases

Non-Metal

Nucleus

Period

Proton

Pure Substance

Stable outer shell

Subatomic particle

Transition metals

Valence electrons

Covalent Compound

Electrons

Element

Family/Group

Halogens

Ionic bonding

Ionic compounds

Ions

Lewis Diagram

Matter

Metal

Metalloids

Mixture

Mechanical

Suspensions

Solutions Elements Compounds

No new substances

produced

Only a change in state or

appearance

New substances produced

Very hard to reverse

= CHEMICAL REACTION

• An atom is the smallest particle of an element that still

has the properties of that element

See pages 168 - 169

50 million atoms, lined

up end to end = 1 cm

An atom = proton(s) +

neutron(s) + electron(s)

• Atoms join together to form compounds.

• A compound is a pure substance that is composed of two or more atoms combined in a specific way.

• Oxygen and hydrogen are atoms/elements; H2O is a compound.

See pages 168 - 169

A chemical change occurs

when the arrangement of

atoms in compounds

changes to form new

compounds.

See pages 168 - 169

• Atoms are made up of smaller particles called subatomic

particles.

See page 170

• The nucleus is at the centre of an

atom.

• The nucleus is composed of

-positive protons

-neutral neutrons

• Electrons exist in the space

surrounding the nucleus.

See page 170

• # of protons = # of electrons in every atom

• Nuclear charge = charge on the nucleus = # of protons

• Nuclear charge = Atomic number

• Atomic number = # of protons = # of electrons

See page 170

INC

REA

SIN

G R

EAC

TIV

ITY

THE PERIODIC TABLE

Where are the following?

• Atomic number

See page 172

INC

REA

SIN

G R

EAC

TIV

ITY

• In the periodic table elements are listed in order by their atomic number. • Metals are on the left

• The transition metals range from group 3 -12

• Non-metals are on the right

• Metalloids form a “staircase” toward the right side.

See page 171

Metals (left of zig zag line)

Physical Properties of Metals: Shiny, good conductors of

heat and electricity, ductile (make wires) and malleable (thin

sheets). Easily lose electrons. Like to join with non-metals.

Corrode (tarnish/rust).

Nonmetals (right of zig zag line)

Physical Properties of Nonmetals: dull appearance, poor

conductor, brittle (breaks easily), not ductile or malleable.

Easily gain electrons. Like to join with metals, but will bond

to other non-metals.

Metalloids (on both sides of zigzag line)

Physical Properties of Metalloids: have properties of both

metals and nonmetals. Solid, shiny or dull, ductile and

malleable, conduct heat and electricity, but not very well.

THE PERIODIC TABLE

Where are the following?

• Metals

• Non-metals

• Transition metals

• Metalloids

See page 172

INC

REA

SIN

G R

EAC

TIV

ITY

• Rows of elements (across) are called periods.

• All elements in a period have their electrons in the same general area around their nucleus.

• Example: period 3 all have 3 electron shells

See page 171 sodium magnesium aluminum

–Columns of elements are called groups, or families. • All elements in a family have similar properties and bond with other elements

in similar ways.

• Group 1 = alkali metals

• Group 2 = alkaline earth metals

• Group 17 = the halogens

• Group 18 = noble gases

See page 171

1 2 17

18

Group 1 = alkali metals

very reactive metals

want to give away 1 electron

ie: lithium, sodium, potassium...

See page 171

1 2 17

18

Group 2 = alkali earth metals

somewhat reactive metals

want to give away 2 electrons

ie: beryllium, magnesium, calcium...

See page 171

1 2 17

18

Group 17 = halogens

very reactive non-metals

want to accept 1 electron

react with alkali metals

ie: fluorine, chlorine, bromine......

See page 171

1 2 17

18

Group 18 = noble gases

STABLE. Very non reactive gaseous non-metals

ie: helium, neon, argon......

See page 171

1 2 17

18

THE PERIODIC TABLE

Where are the following?

• Period

• Group/Family

•Alkali metals

• Alkaline earth metals

• Halogens

• Noble gases

See page 172

INC

REA

SIN

G R

EAC

TIV

ITY

• Atoms gain and lose electrons to form bonds.

• The atoms become electrically charged particles called ions.

See page 173

• Atoms gain and lose electrons to form bonds.

• Metals lose negative electrons & become positive ions.

• Positive ions are called CATIONS.

See page 173

Some metals are MULTIVALENT and can lose a

varying number of electrons.

For example, iron, Fe, loses either two (Fe2+) or

three (Fe3+) electrons

See page 173

• Atoms gain and lose electrons to form bonds.

• Non-metals gain electrons and become negative ions

• Negative ions are called ANIONS

See page 173

Atoms gain and lose electrons in an attempt to be STABLE.

The noble gases are stable because they have FULL outer

shells of electrons. They don’t need to lose or gain any e-s.

Atoms in each period want to have the same number of

electrons in their outer shell (VALENCE ELECTRONS) as

the noble gases on the end of their period.

See page 173

• Bohr diagrams show how many electrons appear in each electron shell around an atom.

• The first electron shell holds 2 electrons

• The second electron shell holds 8 electrons

• The third electron shell holds 8 electrons

• The fourth electron shell holds 18 electrons

See page 174

The noble gas

elements have

full electron

shells and are

very stable.

• Electrons appear in shells in a very predictable manner. • The period number = the number of shells in the atom. • Except for the transition elements (family 3-12), the last digit of the

group number = the number of electrons in the valence shell.

See page 175

See page 174

What element is this?

• It has 2 + 8 + 8 = 18

electrons, and

therefore, 18 protons.

• It has three electron

shells, so it is in

period 3.

• It has eight electrons

in the outer (valence)

shell.

18 p

22 n

argon

• When two atoms get close together, their valence electrons interact.

• If the valence electrons can combine to form a low-energy bond, a

compound is formed.

• Each atom in the compound attempts to have a ‘full’ outer shell of valence

electrons.

See pages 176 - 177

There are 2 types of compounds:

• IONIC COMPOUND: metals lose electrons and non-metals gain

electrons.

• Ionic bonds form when electrons are transferred from positive (+) ions to

negative (-) ions.

• The negative and positive ions are ATTRACTED to each other and form a

BOND.

See pages 176 - 177

• Example ionic bond:

• lithium and oxygen form an ionic bond in the compound Li2O.

See pages 176 - 177

lithium oxygen

+

Electrons are transferred from the positive

ions to negative ions

Li+ O2- Li+

lithium oxide, Li2O

There are 2 types of compounds:

• COVALENT COMPOUND: atoms share electrons.

• Covalent bonds form when electrons are shared between

two non-metals.

• Electrons stay with their atom but overlap with other shells.

See pages 176 - 177

• Example covalent bond

• Hydrogen and fluorine form a covalent bond in the compound

hydrogen fluoride.

See pages 176 - 177

hydrogen fluorine

+

electrons are shared

Hydrogen fluoride

• Lewis diagrams illustrate chemical bonding by

showing only an atom’s valence electrons and

the chemical symbol.

See page 178

Dots representing electrons

are placed around the element

symbols at the points of the

compass (north, east, south,

and west).

• Dots representing electrons are placed around the element symbols at

the points of the compass (north, east, south, and west).

• Electron dots are placed singly until the fifth electron is reached then

they are paired.

See page 178

• Lewis diagrams and IONIC BONDS:

• For positive ions, one electron dot is removed from the valence shell for each positive charge.

• For negative ions, one electron dot is added to each valence shell for each negative charge.

• Square brackets are placed around each ion to indicate transfer of electrons.

See page 179

Be Cl

• •

• •

• •

• •

• •

• •

• •

• •

Each beryllium has two

electrons to transfer away,

and each chlorine can

receive one more electron.

Be Cl Cl

• •

• •

• •

• •

• •

• •

• •

• •

• •

• •

• •

• •

Be Cl Cl

• •

• •

• •

• •

• •

• •

• •

• •

• •

• •

• •

• •

Since Be2+ can donate two

electrons and each Cl– can

accept only one, two Cl– ions

are necessary.

beryllium chloride

2+ – –

• Lewis diagrams and COVALENT BONDS:

–Like Bohr diagrams, valence electrons are drawn to show sharing of electrons.

–The shared pairs of electrons are usually drawn as a straight line.

See page 179

See page 180

• DIATOMIC MOLECULES, like O2 and H2, are also easy to draw as

Lewis diagrams.

The elements Hydrogen, Nitrogen, Fluorine, Oxygen, Iodine, Chlorine,

and Bromine are always found as diatomic molecules.

MEMORY TRICK: I Have No Bright Or Clever Friends