1. General Chemistry Chem 110
Revision
Chapter 5-7-8

2. Chapter 5
5.1 Substances that exist as gases.
5.2 Pressure of gas.
5.3 The gas laws.
5.4 The ideal gas equations.
5.5 Gas Stoichiometry.
5.6 Dalton’ s law of partial pressures.

3. Units of Pressure
1 pascal (Pa) = 1 N/m2
1 atm = 76 cmHg
=760 mmHg = 760 torr
1 atm = 101,325 Pa= x105Pa
1 atm = X102 kPa

4. P1 x V1 = P2 x V2 V1/T1 = V2 /T2 5.3 The gas laws. PV = nRT PM RT d =
Boyle’s Law
Charles`s and Gay-Lussac`s Law
P1 x V1 = P2 x V2
V1/T1 = V2 /T2
5.3 The gas laws.
Ideal Gas Equation
Avogadro`s Law
PV = nRT
V1 / n1 = V2 / n2
General gas law:
Constant moles
The conditions 0 0C and 1 atm are called standard temperature and
pressure (STP).
Experiments show that at STP, 1 mole of an ideal gas occupies L. PM RT
d =
dRT P M=

5. Gas Stoichiometry PT = PO + PH O Dalton’s Law of Partial Pressures
Pi = Xi PT
mole fraction (Xi) = ni nT
PT = PO + PH O 2

6. CHAPTER 7 7.1 From classical physics to quantum theory.
7.3 Bohr’s theory of the hydrogen atom.
7.6 Quantum numbers .
7.7 Atomic orbitals .
7.8 Electron configurations .

7. All electromagnetic radiation
Electromagnetic waves
All electromagnetic radiation
l x n = c
E = h x n = h c / l
= n h n
Planck’s constant (h)
h = 6.63 x J•s
n = 1, 2,3, ……..
1 nm=10-9 m
En = -RH
( ) 1 n2 i f
∆ E = RH
( ) 1 n2

8. 7.6 Quantum numbers . (n2) 7.6 (2l+1) For certain value of (n),
there are (n2) no. of orbitals and (2n2) no. of electrons
For each l there are (2l+1) orbital
7.6

9. Pauli exclusion principle - no two electrons in an atom
can have the same four quantum numbers
The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins (Hund’s rule).
ns2 (n-1)d4 Less stable
ns2 (n-1)d9 less stable
ns1 (n-1)d5 more stable
ns1 (n-1)d10 more stable
7.3

10. CHAPTER 8 8.2 Periodic classification of the elements
8.3 Periodic variation in physical properties
8.4 Ionization energy-
8.5Electron affinitY

11. ns2np6 ns1 ns2np1 ns2np2 ns2np3 ns2np4 ns2np5 ns2 d10 d1 d5 4f 5f 8.2
Ground State Electron Configurations of the Elements (IUPAC)
ns1
ns2np1
ns2np2
ns2np3
ns2np4
ns2np5
ns2
d10 d1 d5 4f 5f
8.2

12. ionic compound that forms from magnesium and oxygen(MgO).
Representing Free Elements in Chemical Equations
Chemists always use the empirical formulas to represent metals and metalloids in chemical equations, such as Fe, Na, and B, Si, Ge.
For nonmetals there is no single rule.
- Carbon, for example, exists in its empirical formula (C) but hydrogen, nitrogen, oxygen, and halogens exist as diatomic molecules [H2, N2, O2,and X2 (F2, Cl2, Br2….].
- The stable form of phosphorus is P4 and S8 for sulfur but for sulfur, chemists often use the empirical formula (S) in chemical equations, rather than S8
3.All the noble gases are monoatomic species.
S + O S O2
ionic compound that forms from magnesium and oxygen(MgO).

13. Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne.
isoelectronic is the atoms have the same number of electrons, and hence the same ground- state electron configuration.
1-Effective nuclear charge(Zeff)
Zeff= Z-σ
2-Atomic radius is one-half the distance between the two nuclei in two adjacent metal atoms (a) or in a diatomic molecule (b).
Decreasing atomic radius
Increasing atomic radius

14. Row 4 O2− < F− < Mg2+ < N3−
3-Ionic radius is the radius of a cation or an anion.
Cation is always smaller than atom from which it is formed.
Anion is always larger than atom from which it is formed.
1- Arrange these ions in order of increasing ionic radius: Mg2+, N3−, O2−, F−   
Increasing radius →
Row 1 Mg2+ < F− < O2− < N3−
Row 2 F− < Mg2+ < O2− < N3−
Row 3 Mg2+ < O2−< N3−< F−
Row 4 O2− < F− < Mg2+ < N3−

15. 4-Ionization energ Increasing First Ionization Energy
Decreasing First Ionization Energy
1- 2A and 3A elements (example ,Be and B
The ionization energies for B(group 3)are lower than those for(Be) group (2A )elements in the same period.
2-group 5A and 6A (example,between N and O and between P and S).The ionization energies for group 6(O) elements are lower than those for group 5A elements (N)in the same period

16. 5-Electron affinity Increasing Electron affinity
Decreasing Electron affinity
1-The electron affinity of a group 2A element is lower than that for the corresponding group 1A element.and electron affinity of a group 5A element(N) is lower than that for the corresponding group 4A(C) element
2-The noble gses have extremely low electron affinities (zero or negative values).
3- The electron affinity of (F) is lower than that for the( Cl)
7-Which of these elements has the greatest electron affinity (largest positive value)?  Mg Al Si P