1. 4.1: Atomic Theory
& BONDING

2. 4.1: Learning Outcomes
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)

3. 4.1: Learning Outcomes
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.

4. 4.1: VOCABULARY Alkali earth metals Alkali metals Anions Atomic #
Atomic number
Atomic Theory
Atoms
Bohr diagram
Cations
Chemical Change
Chemical reaction
Compound
Covalent bonding
Covalent Compound
Electrons
Element
Family/Group
Halogens
Ionic bonding
Ionic compounds
Ions
Lewis Diagram
Matter
Metal
Metalloids
Mixture
Molecule
Neutron
Noble gases
Non-Metal
Nucleus
Period
Proton
Pure Substance
Stable outer shell
Subatomic particle
Transition metals
Valence electrons

5. Chemistry is the study of...
MATTER!!!

6. What is matter? Anything with mass and volume!
Found in 3 phases: liquid, solid, gas
Can't be created or destroyed, it only changes form

7. MATTER Mixtures Pure Substances Compounds Solutions Elements
Mechanical
Suspensions

8. ATOMS
An atom is the smallest particle of an element that still has the properties of that element
50 million atoms, lined up end to end = 1 cm
An atom = proton(s) + neutron(s) + electron(s)
See pages
(c) McGraw Hill Ryerson 2007

9. An atom is the basic structure from which all matter is composed, as a brick is basic to the structure of a wall.

10. Gold is one example of an atom/element.
A bar of gold can be shaved into gold dust, and still be recognizable as gold.
How fine can the dust become and still be considered gold?
The smallest particle that would still have the properties associated with gold is an atom. How small is an atom? Consider that a small gold coin may contain over 20,000,000,000,000,000,000,000 atoms.

11. ATOMS FORM COMPOUNDS 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
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12. ATOMS FORM COMPOUNDS

13. CHEMICAL CHANGE
A chemical change occurs when the arrangement of atoms in compounds changes to form new compounds.
See pages
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14. CHEMICAL CHANGE
Sodium Na (solid)
Chlorine
Cl (gas)
Salt
NaCl

15. ATOMIC THEORY
Atoms are made up of smaller particles called subatomic particles.
See page 170
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16. ATOMIC THEORY
If the proton & neutron were enlarged, and each had the mass of a hippopotamus, the electron, enlarged to the same scale, would have less mass than an owl.

17. ATOMIC THEORY 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
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18. ATOMIC THEORY # 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
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19. INCREASING REACTIVITY

20. The Periodic Table Where are the following? See page 172 Atomic number
INCREASING REACTIVITY
Where are the following?
Atomic number
See page 172
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21. Organization of the Periodic Table
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
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22. 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.

24. The Periodic Table Where are the following? See page 172 Metals
INCREASING REACTIVITY
Where are the following?
Metals
Non-metals
Transition metals
Metalloids
See page 172
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25. Organization of the Periodic Table
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
sodium
magnesium
aluminum
See page 171
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26. Organization of the Periodic Table
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 18 1 2 17
See page 171
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27. Organization of the Periodic Table
Group 1 = alkali metals
very reactive metals
want to give away 1 electron
ie: lithium, sodium, potassium... 18 1 2 17
See page 171
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28. Organization of the Periodic Table
Group 2 = alkaline earth metals
somewhat reactive metals
want to give away 2 electrons
ie: beryllium, magnesium, calcium... 18 1 2 17
See page 171
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29. Organization of the Periodic Table
Group 17 = halogens
very reactive non-metals
want to accept 1 electron
react with alkali metals
ie: fluorine, chlorine, bromine...... 18 1 2 17
See page 171
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30. Organization of the Periodic Table
Group 18 = noble gases
STABLE. Very non reactive gaseous non-metals
ie: helium, neon, argon...... 18 1 2 17
See page 171
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31. The Periodic Table Where are the following? See page 172 Period
INCREASING REACTIVITY
Where are the following?
Period
Group/Family
Alkali metals
Alkaline earth metals
Halogens
Noble gases
See page 172
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33. Periodic Table & Ion Formation
Atoms gain and lose electrons to form bonds.
The atoms become electrically charged particles called ions.
See page 173
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34. Periodic Table & Ion Formation
Atoms gain and lose electrons to form bonds.
Metals lose negative electrons & become positive ions.
Positive ions are called CATIONS.
See page 173
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36. Periodic Table & Ion Formation
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
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37. Periodic Table & Ion Formation
Atoms gain and lose electrons to form bonds.
Non-metals gain electrons and become negative ions
Negative ions are called ANIONS
See page 173
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38. Periodic Table & Ion Formation
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
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39. BOHR MODELS
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
The noble gas elements have full electron shells and are very stable.
See page 174
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40. Patterns of Electron Arrangement in Periods & Groups
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
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41. BOHR MODELS
It has = 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.
What element is this?
18 p
22 n
argon
See page 174
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43. Forming Compounds
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
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44. Forming Compounds 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
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47. Electrons are transferred from the positive ions to negative ions
Forming Compounds
Example ionic bond:
lithium and oxygen form an ionic bond in the compound Li2O. +
lithium
oxygen
Electrons are transferred from the positive ions to negative ions
Li+ O2- Li+
lithium oxide, Li2O
See pages
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48. Forming Compounds 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
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51. Forming Compounds Example covalent bond
Hydrogen and fluorine form a covalent bond in the compound hydrogen fluoride. +
Hydrogen fluoride
hydrogen
fluorine
electrons are shared
(c) McGraw Hill Ryerson 2007
See pages

52. Lewis Diagrams
Lewis diagrams illustrate chemical bonding by showing only an atom’s valence electrons and the chemical symbol.
Dots representing electrons are placed around the element symbols at the points of the compass (north, east, south, and west).
See page 178
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53. Lewis Diagrams
Electron dots are placed singly until the fourth electron is reached then they are paired.
See page 178
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54. Lewis Diagrams

55. Lewis Diagrams To write IONS using lewis diagrams follow these steps:
Step 1: Write the lewis diagram as you normally would.
Step 2: If the element has a POSITIVE combining capacity it will give away an electron and become a POSITIVE ION (cation). Rewrite the lewis diagram to show the element symbol in square brackets (no electrons needed as it has given them away and they now have an EMPTY outer electron shell!) then add the + charge on the outside of the brackets.
Step 3: If the element has a NEGATIVE ION (anion). Rewrite the lewis diagram to show the element symbol in square brackets with extra electrons. They will now have a FULL OUTER electron shell. Then add the - charge on the outside of the brackets.

56. Lewis Diagrams of Ions 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. – 2+ –
• •
• •
• •
• •
• •
• •
• •
• •
• •
• •
• •
• •
• •
• •
• •
• • Be Cl Cl Be Cl Cl Be Cl
Each beryllium has two electrons to transfer away, and each chlorine can receive one more electron.
Since Be2+ can donate two electrons and each Cl– can accept only one, two Cl– ions are necessary.
beryllium chloride
See page 179
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57. Lewis Diagrams of Covalent Bonds
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
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58. Lewis Diagrams of Diatomic Molecules
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
See page 180
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