1.1
Atoms, Electrons, and Orbitals
"Protons"
Protons
positively charged
mass = 1.6726 X 10-27 kg
Neutrons
neutral
mass = 1.6750 X 10-27 kg
Electrons
negatively charged
mass = 9.1096 X 10-31 kg

"Atomic number (Z) =..."
Atomic number (Z) = number of protons in nucleus
(this must also equal the number of electrons
in neutral atom)
Mass number (A) = sum of number of protons
+ neutrons in nucleus

"Schr�dinger combined the idea that..."
Schr�dinger combined the idea that an electron
has wave properties with classical equations
of wave motion to give a wave equation for the
energy of an electron in an atom.
Wave equation (Schr�dinger equation) gives a
series of solutions called wave functions (y ).

"Only certain values of y"
Only certain values of y  are allowed.
Each y  corresponds to a certain energy.
The probability of finding an electron at a
particular point with respect to the nucleus is
given by y 2.
Each energy state corresponds to an orbital.

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"Each orbital is characterized by..."
Each orbital is characterized by a unique
set of quantum numbers.
The principal quantum number n is a whole
number (integer) that specifies the shell and is
related to the energy of the orbital.
The angular momentum quantum number is
usually designated by a letter (s, p, d, f, etc)
and describes the shape of the orbital.

"s Orbitals are spherically symmetric"
s Orbitals are spherically symmetric.
The energy of an s orbital increases with the
number of nodal surfaces it has.
A nodal surface is a region where the probability
of finding an electron is zero.
A 1s orbital has no nodes;  a 2s orbital has one;
a 3s orbital has two, etc.

"No two electrons in the..."
No two electrons in the same atom can have
the same set of four quantum numbers.
Two electrons can occupy the same orbital
only when they have opposite spins.
There is a maximum of two electrons per orbital.

"Principal quantum number (n)..."
Principal quantum number (n) = 1
Hydrogen Helium
Z = 1 Z = 2
1s 1 1s 2

"p Orbitals are shaped like..."
p Orbitals are shaped like dumbells.
Are not possible for n = 1.
Are possible for n = 2 and higher.
There are three p orbitals for each value
of n (when n is greater than 1).

"p Orbitals are shaped like..."
p Orbitals are shaped like dumbells.
Are not possible for n = 1.
Are possible for n = 2 and higher.
There are three p orbitals for each value
of n (when n is greater than 1).

"p Orbitals are shaped like..."
p Orbitals are shaped like dumbells.
Are not possible for n = 1.
Are possible for n = 2 and higher.
There are three p orbitals for each value
of n (when n is greater than 1).

"Principal quantum number (n)..."
Principal quantum number (n) = 2

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1.2
Ionic Bonds
"An ionic bond is the..."
An ionic bond is the force of electrostatic
attraction between oppositely charged ions

"Ionic bonds are common in..."
Ionic bonds are common in inorganic chemistry
but rare in organic chemistry.
Carbon shows less of a tendency to form cations
than metals do, and less of a tendency to form
anions than nonmetals.

1.3
Covalent Bonds
The Lewis Model of Chemical Bonding
In 1916 G. N. Lewis proposed that atoms
combine in order to achieve a more stable
electron configuration.
Maximum stability results when an atom
is isoelectronic with a noble gas.
An electron pair that is shared between
two atoms constitutes a covalent bond.

Covalent Bonding in H2
Sharing the electron pair gives each hydrogen an electron configuration analogous to helium.

Covalent Bonding in F2
Sharing the electron pair gives each fluorine an electron configuration analogous to neon.

The Octet Rule
The octet rule is the most useful in cases involving covalent bonds to C, N, O, and F.

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1.4
Double Bonds and Triple Bonds
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1.5
Polar Covalent Bonds
and Electronegativity
Electronegativity is a measure of an element to attract electrons toward itself when bonded to
another element.
An electronegative element attracts electrons.
An electropositive element releases electrons.

"Electronegativity increases from left to..."
Electronegativity increases from left to right
in the periodic table.
Electronegativity decreases going down a group.

Generalization
The greater the difference in electronegativity
between two bonded atoms;  the more polar the
bond.

Generalization
The greater the difference in electronegativity
between two bonded atoms;  the more polar the
bond.

1.6
Formal Charge
Formal charge is the charge calculated for an atom in a Lewis structure on the basis of an equal sharing of bonded electron pairs.

Nitric acid
We will calculate the formal charge for each atom in this Lewis structure.

Nitric acid
Hydrogen shares 2 electrons with oxygen.
Assign 1 electron to H and 1 to O.
A neutral hydrogen atom has 1 electron.
Therefore, the formal charge of H in nitric acid is 0.

Nitric acid
Oxygen has 4 electrons in covalent bonds.
Assign 2 of these 4 electrons to O.
Oxygen has 2 unshared pairs.  Assign all 4 of these electrons to O.
Therefore, the total number of electrons assigned to O is 2 + 4 = 6.

Nitric acid
Electron count of O is 6.
A neutral oxygen has 6 electrons.
Therefore, the formal charge of O is 0.

Nitric acid
Electron count of O is 6 (4 electrons from unshared pairs + half of 4 bonded electrons).
A neutral oxygen has 6 electrons.
Therefore, the formal charge of O is 0.

Nitric acid
Electron count of O is 7 (6 electrons from unshared pairs + half of 2 bonded electrons).
A neutral oxygen has 6 electrons.
Therefore, the formal charge of O is -1.

Nitric acid
Electron count of N is 4 (half of 8 electrons in covalent bonds).
A neutral nitrogen has 5 electrons.
Therefore, the formal charge of N is +1.

Nitric acid
A Lewis structure is not complete unless formal charges (if any) are shown.

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