Mar 19, 2016
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., O‐ 2, 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 metalsAlkali metalsAnionsAtomic #Atomic numberAtomic TheoryAtoms Bohr diagramCations Chemical ChangeChemical reactionCompound Covalent bonding
MoleculeNeutronNoble gasesNon-MetalNucleusPeriodProtonPure SubstanceStable outer shellSubatomic particleTransition metalsValence electrons
Covalent CompoundElectronsElement Family/GroupHalogensIonic bondingIonic compoundsIonsLewis DiagramMatterMetalMetalloidsMixture
MechanicalMechanicalSuspensionsSuspensions
SolutionsSolutions ElementsElements CompoundsCompounds
(c) McGraw Hill Ryerson 2007
• 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)
(c) McGraw Hill Ryerson 2007
• 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
(c) McGraw Hill Ryerson 2007
A chemical change occurs when the arrangement of atoms in compounds changes to form new compounds.
See pages 168 - 169
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• Atoms are made up of smaller particles called subatomic particles.
See page 170
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.
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• 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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– # 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
INCR
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(c) McGraw Hill Ryerson 2007
The Periodic Table
Where are the following?• Atomic
number
See page 172
INCR
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REAC
TIVI
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(c) McGraw Hill Ryerson 2007
• 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.
(c) McGraw Hill Ryerson 2007
The Periodic Table
Where are the following?• Metals
• Non-metals
• Transition metals
• Metalloids
See page 172
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– 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
(c) McGraw Hill Ryerson 2007
–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 1718
(c) McGraw Hill Ryerson 2007
Group 1 = alkali metalsvery reactive metalswant to give away 1 electronie: lithium, sodium, potassium...
See page 171
1 2 17
18
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Group 2 = alkaline earth metalssomewhat reactive metalswant to give away 2 electronsie: beryllium, magnesium, calcium...
See page 171
1 2 17
18
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Group 17 = halogensvery reactive non-metalswant to accept 1 electronreact with alkali metalsie: fluorine, chlorine, bromine......
See page 171
1 2 17
18
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Group 18 = noble gasesSTABLE. Very non reactive gaseous non-metalsie: helium, neon, argon......
See page 171
1 2 17
18
(c) McGraw Hill Ryerson 2007
The Periodic Table
Where are the following?• Period
• Group/Family
•Alkali metals
• Alkaline earth metals
• Halogens
• Noble gases
See page 172
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• Atoms gain and lose electrons to form bonds.– The atoms become electrically charged
particles called ions.
See page 173
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• 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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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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• 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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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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• 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.
(c) McGraw Hill Ryerson 2007
• 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
(c) McGraw Hill Ryerson 2007 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
(c) McGraw Hill Ryerson 2007
• 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
(c) McGraw Hill Ryerson 2007 See pages 176 - 177
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• 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
(c) McGraw Hill Ryerson 2007
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
(c) McGraw Hill Ryerson 2007
• 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
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• 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).
(c) McGraw Hill Ryerson 2007
–Electron dots are placed singly until the fourth electron is reached then they are paired.
See page 178
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.
(c) McGraw Hill Ryerson 2007
• 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.
BeCl Cl• •
• •• •
• •
• •• •• •
• •
• •• •• •
• •
BeCl Cl• •
• •• •
• •
• •• •• •
• •
• •• •• •
• •Since Be2+ can donate two electrons and each Cl– can
accept only one, two Cl– ions are necessary.
beryllium chloride
2+ ––
(c) McGraw Hill Ryerson 2007
• 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
(c) McGraw Hill Ryerson 2007 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.
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