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General Chemistry (Chem110)
Dr. Rima Alharthy
Assistant professor in Medicinal and organic synthesis
University of Nottingham, UK
[email protected]
Building 4 room 204
Office hours (Mondays and Wednesdays 8:00-9:30 a.m.)
Chapter 2 (p1-p31)
Chapter 2: Atoms, molecules and ions
2.1-The Atomic theory
2.3-Atomic number, mass number and isotopes
2.4-The periodic table
2.5-Molecules and ions
2.6-Chemical formulas
2.7-Naming compounds
Chapter 2: Dalton’s Atomic Theory (1808)
•
Elements
are composed of extremely small particles called atoms.
Examples: Na, Cl, Ca and H2
•
All atoms of a given element are identical, having the same size,
mass and chemical properties.
•
The atoms of one element are different from the atoms of all
other elements.
Chapter 2: Dalton’s Atomic Theory (1808)
•
Compounds
are composed of atoms of more than one element.
Examples: H2O, NaCl and CaCO3
•
Chemical reaction
involves only the separation, combination, or rearrangement of
atoms; it does not result in their creation or destruction.
Chapter 2: The structure of the atom
Neutron ±
Proton +
Electron • Protons and electrons are the only particles that have a charge.
• Protons and neutrons have essentially the same mass.
• The mass of an electron is so small we ignore it.
Chapter 2: Atomic number and mass number
Atomic number (Z)
Number of protons in nucleus
Mass number (A) = Number of protons + number of neutrons
= atomic number (Z) + number of neutrons
Atomic number
Mass number
11
Na
22.98
Sodium
Element symbol
Element name
Example 2.1
Give the number of protons, neutrons and electrons in each of the
following species:
Element
Mass Number
Atomic Number
Number of electrons
Number of protons
Number of neutrons
20
11
Na
22
11
Na
17
8
O
14
6
C
Chapter 2: isotopes
Isotopes
Atoms that have the same atomic number (i.e protons and electrons)
but different mass numbers (i.e. neutrons)
Examples
Active, used in
nuclear weapons
235
92
U
238
92
U
Not active
Chapter 2: isotopes
Some applications of isotopes
•
Nuclear reactors and weapons.
•
Radiation oncology.
•
Spectroscopy.
•
Chemical labelling technique.
Properties of isotopes
•
The chemical properties of an element are determined primarily
by the protons and electrons.
•
Neutrons do not take a part in chemical changes under normal
conditions.
Chapter 2: Periodic table
Alkaline
Metal
Group
1A
Alkaline
Earth
Metal
Group
2A
Main group elements
1A-7A
Group
Halides
7A
Noble
Gases
8A
Transition metals
1-8B
Period
Lanthanides
Actanides
Chapter 2: molecules
Molecule
An aggregate of two or more atoms in a definite arrangement held
together by chemical forces (bonds).
Diatomic molecule
polyatomic molecule
contains only two atoms
contains more than two atoms
H2, N2, O2, Br2, HCl, CO
O3, H2O, NH3, CH4
Chapter 2: ions
Ion
is an atom, or group of atoms, that has a net positive or negative charge.
Cation
Na
- 1 electron
11 electrons
Anion
+ 1 electron
Na+
Cl
10 electrons
17 electrons
If a neutral atom loses one or
more electrons it becomes a cation
has a positive charge
Cl-
18 electrons
If a neutral atom gains one or more
electrons it becomes an anion.
Anion has a negative charge
Chapter 2: ions
Cation
Monoatomic
cation
Polyatomic
cation
Na+, K+, Mg2+
NH4+
Anion
Monoatomic
anion
Cl-, Br-,S2-
Polyatomic
anion
OH-
Chapter 2: Chemical formulas
Molecular formula
Shows the exact number of atoms of each element in the smallest unit
of a substance
Empirical formula
Shows the simplest whole-number ratio of the atoms in a substance
Molecular
H 2O
C6H12O6
empirical
H 2O
CH2O
O3
O
N 2H 4
NH2
Chapter 2 Formula of ionic compounds
Ionic compounds
Consist of a combination of cations (metal) and anions (non metal)
The sum of the charges on the cation(s) and anion(s) in each formula
unit must equal zero
The ionic compound NaCl
Chapter 2: Ionic formula
Example 2.4
Write the formula of
a) magnesium nitride containing Mg2+ and N3- ions.
b) Chromium sulfate containing Cr3+ and SO42- ions.
c) Titanium oxide containing Ti4+ and O2- ions.
Chapter 2 Naming compounds
Chemical nomenclature of inorganic compounds
1- Ionic compounds
2- molecular compounds
3- acids and bases
4- hydrates
Chapter 2: Naming ionic compounds
Ionic Compounds
Formed of metal + nonmetal
Rules of naming ionic compounds
• Starts with the cation (metal) original name.
• End with anion (nonmetal), add -ide to element name
Examples
Mg(OH)2
magnesium hydroxide
KNO3
potassium nitrate
BaCl2
barium chloride
K 2O
potassium oxide
Chapter 2: Naming ionic compounds
Cation
Anion
Aluminium Al3+
Ammonium NH4+
Barium Ba2+
Cadmium Cd 2+
Calcium Ca2+
Cesium Cs+
Chromium(III)Cr3+
Cobalt (II) Co2+
Copper (I) Cu+
Copper (II) Cu2+
Hydrogen H+ (Proton)
Iron (II) Fe2+
Iron (III) Fe3+
Lead(II) Pb2+
Lithium Li+
Magnesium Mg2+
Manganese (II) Mn2+
*Mercury (I) Hg22+
Mercury (II) Hg2+
Potassium K+
Rubidium Rb+
Silver Ag+
Sodium Na+
Strontium Sr2+
Tin (II) Sn2+
Tin (IV) Sn4+
Zinc Zn2+
Bromide Br Carbonate (CO3)2Chlorate (ClO3)Chloride ClChromate (CrO4)2Cyanide CNDichromate (Cr2O7)2Dihydrogen phosphate (H2PO4)Fluoride FHydride HHydrogen carbonate or bicarbonate (HCO3)Hydrogen phosphate (HPO4)2Hydrogen sulfate or bisulfate (HSO4)Hydroxide OHIodide INitrate (NO3)Nitride NNitrite (NO2)Oxide O2Permenganate (MnO4)Peroxide O22Phosphate (PO4)3Sulfate (SO4)2Sulfide S2Sulfite (SO3)2Thiocyanate (SCN)-
Chapter 2: Naming ionic compounds
• Transition metal ionic compounds
can form more than one cation. Therefore, old Roman nomenclature
system (Stock system) applies.
Examples
FeCl2
Iron (II) chloride
FeCl3
Iron (III) chloride
Cr2S3
Chromium (III) sulfide
Chapter 2: Example 2.5
Name the following compounds:
a) Cu(NO3)2
b) KH2PO4
c) NH4ClO3
d) PbO
e) Li2SO3
Chapter 2: Example 2.6
Write chemical formulas for the following compounds
a) Mercury (I) nitrite
b) Cesium sulfide
c) Calcium phosphate
d) Rubidium sulfate
e) Barium hydride
Chapter 2: naming molecular compounds
Molecular binary compounds
Formed of nonmetal + nonmetal
or
nonmetals + metalloids
Examples
H2O, NH3 and CH4
Rules of naming molecular compounds
• Starts with element nonmetal
• End with –ide for nonmetal
• Omit prefix (eg. mono for first element.
• For oxides omit the ending ‘a’ and ‘o’ in the prefix.
Chapter 2: naming molecular compounds
Examples
HI
hydrogen iodide
SO2
sulfur dioxide
N2Cl4
dinitrogen tetrachloride
NO2
nitrogen dioxide
N 2O
dinitrogen monoxide
Chapter 2: Example 2.7
Name the following molecular compounds:
a) SiCl4
b) P4O10
c) NF3
d) Cl2O7
Chapter 2: Example 2.8
Write chemical formulas for the following compounds
a) Carbon disulfide
b) Disilicon hexabromide
c) Sulfur tetraflouride
d) Dinitrogen pentoxide
Chapter 2: naming acids and bases
Acid
a substance that yields hydrogen ions (H+) when dissolved in water.
Example HCl
•Pure substance
•Dissolved in water
hydrogen chloride
hydrochloric acid
Chapter 2: naming acids and bases
Rules of naming acids
• Starts with Hydro
• Ends with element + ic + acid
Chapter 2: Naming bases
Base
a substance that yields hydroxide ions (OH-) when dissolved in water.
Rules of naming basse
• Starts with element
• Ends with hydroxide
Examples
NaOH
Sodium hydroxide
KOH
Potassium hydroxide
Ba(OH)2
Barium hydroxide