Download Atoms, Molecules and Ions Chapter 2

Atoms, Molecules and Ions
Chapter 2
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Dalton - 1808
• Atoms of different elements could be distinguished by
differences in their weights.
–
–
–
–
–
–
All matter is composed of atoms.
Atoms cannot be made or destroyed.
All atoms of the same element are identical.
Different elements have different types of atoms.
Chemical reactions occur when atoms are rearranged.
Compounds are formed from atoms of the constituent elements.
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Thomson - 1897
• Discovered the electron.
• Calculated charge to mass ratio
• Plum pudding model
Millikan - 1909
•
•
•
•
Oil Drop Experiment
Measured the charge and mass of the electron
Charge: 1.602 x 10-19 C
Mass: 9.10 x 10-28 g
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Rutherford -1911 & Chadwick 1932
• Experimental work with α- particles
• The positive charge must be concentrated in a tiny
volume at the centre of the atom, while the electrons
orbited.
• 1919 – Discovered positive
particles (protons).
• Chadwick - neutrons
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Atoms
• Proton
+ Electron
+ Neutron
(1.602 x 10-19 C)
(-1.602 x 10-19 C)
+1
-1
• Atomic mass is small so we use atomic mass
units (amu) represented by u.
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Atomic Mass
Charge
Proton
Electron
Neutron
Mass
Actual in C Relative
1.602 x 10-19
+1
Actual in g
1.673 x 10-24
Relative in u
1.00727
1.602 x 10-19
0
9.109 x 10-28
1.675 x 10-24
0.00054858
1.00866
-1
0
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↔
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No of atoms in a teaspoon of salt
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Differentiating elements
• Each element has a characteristic number of PROTONS –
called atomic number (Z).
• The number of protons = no. electrons
• No of neutrons can be different for an element.
• The number of protons + neutrons = mass number (A).
• Atoms with differing numbers of neutrons are called
isotopes.
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Mass number =
Atomic number =
Mass number =
protons + neutrons
No of protons =
No. of electrons =
No of neutrons =
Mass number =
Atomic number =
Atomic number =
no. of protons/ electrons
No of protons =
No. of electrons =
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No of neutrons =
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Isotope examples…
Isotope
Mass
number
Hydrogen -1
1H
Hydrogen -2
2H
Hydrogen -3
3H
Protons
Electrons
Neutrons
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Cobalt
Isotope
Mass
number
Cobalt -56
56Co
59
Cobalt -57
57Co
60
Cobalt -58
58Co
61
Protons
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Electrons
Neutrons
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Atomic Mass
• Mass of an atom is measured relative to carbon
• 12 u = 12C
• e.g. 1H = 1.008 u (12 x lighter than 12C)
16O = 15.9949 u
• Since most elements occur as a mixture of isotope –
use AVERAGE atomic mass
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Average atomic mass
e.g. Calculate the average atomic mass of Ag
Isotope
107Ag
109Ag
Mass / u
106.90509
108.90476
% Abundance
51.84
48.16
From 107Ag =
From 109Ag =
Atomic mass =
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Try
1. 52.0 % of Br has 44 neutrons= mass 79.0 [
48.0 % of Br has 46 neutrons = mass 81.0 [
Calculate the atomic mass of Br.
]
]
2. Boron (Z = 5) has two naturally occurring isotopes.
Find the % abundance of 10B and 11B given the
following:
10B
11B
Isotopic mass
10.129 u
11.0093 u
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Atomic mass
10.81 u
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The periodic table
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The periodic table
Atomic number
Atomic symbol
Atomic weight
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Chemical Compounds
Chapter 3
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Molecules and Molecular Compounds
• A molecule consists of two or more atoms bound
together
• Each molecule has a chemical formula which indicates
– Which atoms are found in the molecule, and
– In what proportion they are found.
• A molecule made up of two of the same atoms is called
a diatomic molecule. (Halogens – e.g. Cl, Br)
• Compounds composed of molecules are molecular
compounds.
• Most molecular substances contain only nonmetals.
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Examples
H2
H–H
diatomic
H2O
H–O–H
NH3
polyatomic
CH4
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Molecular and Empirical Formulas
• Molecular formulas: Give the actual numbers and
types of atoms in a molecule.
• Empirical formulas : Give the relative numbers and
types of atoms in a molecule (they give the lowest
whole-number ratio of atoms in a molecule).
Molecular Formula
Empirical Formula
H2O
H2O
H2O2
HO
C2H4
CH2
C6H12O6
CH2O
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Ions and Ionic Compounds
• If electrons are added to or removed from a neutral
atom, an ion is formed.
• When an atom or molecule loses electrons it becomes
positively charged.
– Positively charged ions are called cations (M+)
11 e11 p+
Na atom
10 e11 p+
Na+ ion
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When an atom or molecule gains electrons it becomes
negatively charged.
– Negatively charged ions are called anions (M-).
17 e-
18 e-
17 p+
17 p+
Cl atom
Cl- ion
• Generally, metal atoms tend to lose electrons and
non-metal atoms gain electrons.
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Predicting Ionic Charges
• An atom or molecule can gain or lose more than one
electron.
• Many atoms gain or lose enough electrons to have
the same number of electrons as the nearest noble
gas (group 8A).
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Ionic Charges
+3 ±4 -3
+1 +2
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-2 -1
0
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Examples
•
•
•
•
P
Al
Ca
Si
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Oxidation state
• Used to determine the ratio in which the elements forming a
compound will exist.
• Use the ionic charge.
• Follow a number of rules
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Oxidation State
1. Free elements are assigned an oxidation state of 0.
2. The oxidation state for any simple one-atom ion is equal to its
charge.
3. The alkali metals (Li, Na, K, Rb, Cs and Fr) in compounds are
always assigned an oxidation state of +1.
4. Fluorine in compounds is always assigned an oxidation state
of -1.
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5. The alkaline earth metals (Be, Mg, Ca, Sr, Ba, and Ra) and
also Zn and Cd in compounds are always assigned an
oxidation state of +2. Similarly, Al & Ga are always +3.
6. Hydrogen in compounds is assigned an oxidation state of
+1.
7. Oxygen in compounds is assigned an oxidation state of -2.
Exception - Peroxide, e.g. H2O2 (O = -1).
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8. The sum of the oxidation states of all the atoms in a species
must be equal to net charge on the species.
e.g. Net Charge of CrO42- = -2,
For Cr: x + (- 4)x 2 = -2, →x = +6, so the oxidation state of Cr is +6.
• K2MnO4
K = +1
+ 2 contribution
O = -2
Mn = [2 x (+2)] + [ 4 x (-2)] = +6
• B4O72B=?
O = -2
-14 contribution
BOS + (-14) = -2
B4 = 12 and each B = +3
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-8 contribution
Net charge = -2
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Try
Determine the oxidation state of
1. P in PCl5
2. Cl in Ca(ClO4)2
3. S in S2O324. S in Na2SO3
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Charges for common polyatomic anions
• You will be taught HOW to determine these
charges in a few weeks.
Anion
OH
NO3
2SO4
CO32PO42-
Charge
-1
-1
-2
-2
-2
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Ionic compounds
• Form when cations and anions attract each
other to form a neutral compound.
• Ratio of ions can differ.
• An ionic compound must have a metal
present.
– H classified as nonmetal.
e.g. Na+ + Cl- → NaCl (+1 from Na and -1 from Cl = 0
net charge)
e.g. Al3+ and O2- → Al2O3 ( 2x +3 from Al and 3 x -2 from
O = 0 net charge
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Al2O3 …
• What does this mean?
• 2 Al3+ ions and 3 O2- ions make up the molecule.
Try :
Ca2+ and NO3Ca: 2+
NO3: -1
1 Ca + 2 NO3 = +2 + (-1)2 = 0
Ca(NO3)2
() implies the subscript applies to ALL the atoms within.
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Molecular
Ionic
H + non-metal
Metal + non-metal
HCl, C2H4
CaCl2, LiF
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Ionic bonds (transferring electrons)
Covalent bonds (sharing electrons)
Compounds
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Naming inorganic ionic compounds
Cations
a) Cations formed from metal ions have the
same name as the metal
b) The positive charge of the metal is indicated
by a roman numeral in brackets.
e.g. Arsenic can exist with a +3 charge and a
+ 5 charge. This is represented as
As(III) and As(IV)
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c) Cations formed from non-metal atoms have
names that end in –ium
e.g. NH4+
ammonium ion
H3O+
hydronium ion
Anions
a) Monoatomic ions have names ending in –ide
e.g. I
Iodine
iodide
BrBromine bromide
** OHhydroxide
CNcyanide
O22peroxide
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b) Polyatomic anions with O have names ending
in
–ate
or
–ite
Same charge but one less O
Most common oxy-anion
SO42- Sulfate ion
SO32- Sulfite ion
c) When an H+ is added to an oxyanion add
hydrogen (1H) or dihydrogen (2H) to the
name.
e.g. H2SO4
dihydrogen sulfate
H2O
dihydrogen oxide
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** Some compounds can have multiple oxyanion forms
ClO-
ClO2-
ClO3-
ClO4-
The number of O's relative to the Chlorine is changing,
but the ionic charge is not (-1 for all).
The -ite and -ate suffixes are still used, but to
distinguish between the four forms we use prefixes:
ClOClO2ClO3ClO4-
hypochlorite ion
chlorite ion
chlorate ion
perchlorate ion
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e.g.
CuBr2
NH4ClO3
KHSO4
NaClO2
Lithium sulfite
Barium phosphate
Try
1. Sodium persulphate
2. Magnesium(II) bromide
3. Potassium cyanide
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Names and formulas of Acids
Acids
• An acid is a substance whose molecules yield
hydrogen (H+) ions when dissolved in water.
• Acid = anion + H+
• The name of the acid is related to the name of
the anion
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• Anions whose names end in -ide have associated acids that
have the hydro- prefix and an -ic suffix:
Cl- chloride anion
HCl
hydrochloric acid
S2- sulfide anion
H2S
hydrosulfuric acid
• For naming the acids of oxyanions:
– If the anion has an -ate ending, the corresponding acid is given an -ic
ending.
– If the anion has an -ite ending, the corresponding acid has an -ous
ending.
• Prefixes in the name of the anion are kept in naming the acid
ClO- hypochlorite ion
HClO hypochlorous acid
ClO2- chlorite ion
HClO2 chlorous acid
ClO3- chlorate ion
HClO3 chloric acid
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ClO4- perchlorate ion
HClO4 perchloric acid
Rules for Binary Molecular Compounds
• The naming system is for compounds
composed of two non-metallic elements.
• The element furthest left is named first.
(except when O and a halogen!)
Prefixes
• The first element gets a prefix if it has a
1 – mono
subscript in the formula.
– The second element gets the –ide
suffix (ending)
– The second element ALWAYS
gets a prefix
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2 – di
3 – tri
4 – tetra
5 – penta
6 – hexa
7 - hepta44
Examples
1. N2O4
2. P4O10
3. PBr5
4. nitrogen trichloride
5. disulfur hexafluoride
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Prefixes
1 – mono
2 – di
3 – tri
4 – tetra
5 – penta
6 – hexa
7 - hepta
Simple Organic Compounds
Alkanes
• Contain only C and H
• Name according to number of carbons, end in –ane.
• After 4 carbons names as for other compounds.
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Alcohols
• When an H is replaced by –OH group.
• -ol added of end of name.
Methanol
ethanol
1-propanol
1- butanol
1-pentanol
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You should be able to …
• understand the concept of the atom.
• differentiate elements according to mass number.
• use the periodic table to determine the number of protons, electrons and
neutrons an element has.
• explain what an isotope is.
• calculate the average atomic mass of an element.
• identify basic groups in the periodic table.
• understand and differentiate between ions, molecules and their
compounds.
• predict the charge for ions.
• Assign oxidation states to elements in a charged or neutral compound
• Determine the name or the chemical formula for ionic and binary molecular
compounds, acids and simple alkanes.
KNOW
• The charges for common polyatomic anions.
• The rules for naming ionic and binary molecular compounds, acids and
simple alkanes.
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