Download General Chemistry

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TOPICS ON CHEMISTRY FOR ADMISSION TEST
GENERAL CHEMISTRY
The SI system of measurement.
Classification of matter.
Elements, symbols of the elements.
The structure of atoms.
Atomic, molecular and molar mass relationships.
Chemical equations and stoichiometry.
Electronic structure of the elements: quantum numbers, orbitals, electron configuration.
Periodic table: main groups, periodic properties.
Types of chemical bonds.
Covalent bond: Lewis structures, molecular shapes, valence bond theory.
Intermolecular forces.
The gaseous state: The gas laws. Stoichiometric relationships with gases. Kinetic - molecular theory of gases.
Liquid and solid states. Phase changes.
Solutions and their properties: Concentration of solutions. Ions in aqueous solution:
electrolytes and nonelectrolytes.
Chemical equilibrium. The equilibrium constant.
Acids and bases: The pH in solutions of strong acids and strong bases. Equilibria in solutions of weak acids
and weak bases.
Thermochemistry: Energy changes and energy conservation. Expansion work. Energy and enthalpy. Hess’s
law.
Oxidation and reduction: Oxidation state. The activity series of the elements. Balancing redox reactions.
Textbook: McMurry, J., Fay, R.C. (2012): Chemistry, 6th Edition. Pearson Education,
Inc., Upper Saddle River, NJ 07458.
The answers of the above-listed topics:
Answers
Atom: is the smallest particle of an element that retains the chemical properties of that element.
The Modern View of Atomic Structure
The atom consists of positive, negative, and neutral entities (protons, electrons, and neutrons).
Protons and neutrons are located in the nucleus of the atom, which is small. Most of the mass of the atom is
due to the nucleus.
There can be a variable number of neutrons for the same number of protons. Isotopes have the same number
of protons but different numbers of neutrons.
Electrons are located outside of the nucleus. Most of the volume of the atom is due to electrons.
One electron mass is 9.109 X 10-28g, proton mass and neutron mass is 1836 times greater than electron
Isotopes, Atomic Numbers, and Mass Numbers
Atomic number (Z) = number of protons in the nucleus. (number of protons = number of electrons)
Mass number (A) = total number of nucleons (i.e., protons and neutrons) in the nucleus.
Isotopes: Atoms have the same Z but different in A example C (carbon) has two main isotopes C12 and C13,
which differ in
one Neutron, C13 has one more neutrons, ie. 7 neutrons instead of 6.
The Atomic Mass Scale
1H weighs 1.6735 x 1024 g.
We define: mass of 12C = exactly 12 amu.
Using atomic mass units:
1 amu = 1.66054 x 1024 g
1 g = 6.02214 x 1023 amu
Formula and Molecular Mass
Formula mass (FM): sum of atomic masses of all atoms in a formula unit of any compound, molecular or
ionic.
FM (H2SO4) = 2AM(H) + AM(S) + 4AM(O)
= 2(1.0 amu) + (32.0 amu) + 4(16.0 amu) = 98.0 amu
Molecular mass: sum of atomic masses of all atoms in a molecule.
MW for (C6H12O6) = 6(12.0 amu) + 12(1.0 amu) + 6(16.0 amu)= 180g
Mole: Convenient measure chemical quantities.
1 mole of something = 6.0221367 x 1023 of that thing.
Experimentally, 1 mole of 12C has a mass of 12 g.
A mole of anything contains the same number of particles as there are carbon atoms in 12.01 g of carbon.
1 mol C = 6.02 x 1023 C atoms
The number 6.02 x 1023 is also known as Avogadro’s number.
Thus, in one mole of any element, there is Avogadro’s number of atoms.
1 mol Na = 6.02 x 1023 Na atoms
1 mol Au = 6.02 x 1023 Au atoms
Molar Mass
Molar mass: the mass of one mole of any substance in grams (units g/mol).
Mass of 1 mole of 12C = 12 g.
Molar mass are numerically equal to the Formula weights.
Examples:
One mole of water contains 6.02 x 1023 molecules of water which weighs 18.02 grams.
One mole of iron metal Fe contains 6.02 x 1023 atoms which weighs 55.85 grams.
Elements: One type of atom. Example Au, Na, Fe, etc...
Compound: Two or more types of atoms combined chemically example: H2O, CO2, NaClO4, etc...
When an atom or molecule loses electrons, it becomes positively charged.
For example, when Na loses an electron it becomes Na+. Positively charged ions are called cations.
When an atom or molecule gains electrons, it becomes negatively charged.
For example when Cl gains an electron it becomes Cl-. Negatively charged ions are called anions.
An atom or molecule can lose more than one electron!
Predicting Ionic Charge
The number of electrons an atom loses is related to its position on the periodic table. Metals tend to form
cations whereas non-metals tend to form anions
Ionic Compounds
The majority of chemistry involves the transfer of electrons between species.
Example:
To form NaCl, the neutral sodium atom, Na, must lose an electron to become a cation: Na+.
The electron cannot be lost entirely, so it is transferred to a chlorine atom, Cl, which then becomes an anion:
Cl-.
The Na+ and Cl- ions are attracted to form an ionic NaCl lattice which crystallizes.
Chemical bonds
We consider three bonds within molecules (intramolecular force):
• ionic bond (electrostatic forces which hold ions together, e.g. NaCl);
• covalent bond (results from sharing electrons between atoms, e.g. Cl2);
• metallic bonding (refers to metal nuclei floating in a sea of electrons, e.g. Na).
Valence electrons are involved in different chemical bonds
VALENCE ELECTRONS: The electrons in the outermost shell of the atom are called VALENCE
ELECTRONS.
For the main group elements, the group number is the number of valence electrons.
OCTET RULE: Elements want to achieve eight valence electrons (Nobel gases)
How can elements achieve an octet by Bonding either:
1- TRANSFER electrons to form IONIC bond.
2- SHARE electrons to form COVALENT bond
Ionic bonding
An ionic bond is a chemical bond formed by the electrostatic attraction between positive and negative ions.
This type of bond involves the transfer of electrons from one atom (usually a metal) to another (usually a
nonmetal).
The number of electrons lost or gained by an atom is determined by its need to be “isoelectronic” with a noble
gas.
Cl
+ e-
Na
ClNa+
Cl-Na+ + e-
Describing Ionic Bonds
Consider the transfer of valence electrons from a sodium atom to a chlorine atom.
The resulting ions are electrostatically attracted to one another: Na + Cl → NaCl
The attraction of these oppositely charged ions for one another is the ionic bond.
Complete transfer of electrons between atoms converts them to ions, and they will form ionic compounds
Ionic compounds tend to form generally between metals (Sodium, Magnesium, Zinc..) and nonmetals(Oxygen, Nitrogen, Chlorine, Fluorine..), transfer electron(s) from metal to non metal.
Ionic compounds form extensive arrays called crystals(Ionic solid)
Sodium and chloride are held together in the crystal by ionic bonds (strong electrostatic interactions)
POSITIVE IONS:
Metals (left side of periodic table) will lose electrons to form positive ions. Positive ions are called
CATIONS.
A (monatomic) metal ion (cation) is named by its element name.
NEGATIVE IONS:
Nonmetals (right side of periodic table) will gain electrons to form negative ions. Negative ions are called
ANIONS.
A (monatomic) anion is named by placing -ide at the end of the root of the element’s name.
Covalent Bonds
Covalent Bonds: the most common kind of chemical bond
When two or more nonmetals bond, they often share electrons since they have similar attractions for them.
This sharing of valence electrons is called the covalent bond.( H2, CH4 , NH3)
Hydrogen, Oxygen, Nitrogen and Halogens are exist a molecules (diatom) rather than one atom (H2,O2, N2
,F2, Cl2, Br2, I2) These atoms will share sufficient numbers of electrons in order to achieve a noble gas
electron configuration (that is, eight valence electrons). The tendency of atoms in a molecule to have eight
electrons in their outer shell (two for hydrogen) is called the octet rule.
O O
double
bond
N
N
Triple Bond
When similar atoms bond, they share pairs of electrons to each obtain an octet (O2, I2, F2 , H2). Each pair of
shared electrons constitutes one chemical bond. Example: H + H  H2 has electrons on a line connecting the
two H nuclei. (O2, I2, F2, CO2, CH4).
Multiple Covalent Bonds
It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds):
Single covalent bond: One shared pair of electrons = single bond (e.g. H2); H-H
Double covalent bond: Two shared pairs of electrons = double bond (e.g. O2); O=O
Triple covalent bond: Three shared pairs of electrons = triple bond (e.g. N2).
Generally, bond distances decrease as we move from single through double to triple bonds.
• In a covalent bond, electrons are shared.
• The shared pair electrons(one, two or three pairs) are located between the nuclei
• Optimum distance between nuclei called the bond strength which corresponds to minimum energy
need to have the most stable molecule.
• Amount of energy released when the bond forms is called bond dissociation energy(D)
• Sharing of electrons to form a covalent bond does not imply equal sharing of those electrons. There
are some covalent bonds in which the electrons are located closer to one atom than the other.
• Unequal sharing of electrons results in polar bonds, example CH4, H2O, HCl
The absolute value of the difference in electronegativity of two bonded atoms gives a rough measure of the
polarity of the bond as follows:
1- When this difference is small (less than 0.5), the bond is nonpolar.
2- When this difference is large (greater than 0.5), the bond is considered polar.
3- When this difference exceeds approximately 1.8, sharing of electrons is no longer possible and
the bond becomes ionic.
Metallic bonding
Is the attraction between positive metal ions and surrounding freely mobile electrons in the outermost shell
(valence electrons). The valence electrons from all atoms formed what has been called “electron sea”
Most substance that have metallic bonding are either metallic elements or alloys (mixture of metals, or non
metals and metals)
Metals with such bonding have high boiling pints because it is difficult to liberate metal ions from the
surrounding free electrons.
Metals are luster because the free electrons can absorb and reemit visible light
Metals are Malleable (can’t be broken easily) and have ductility (drawn out into wires).
Because of the free mobile electrons metals are good conductors for electricity and heat.
The Intermolecular forces:
• The covalent bond holding a molecule together is an intramolecular force.
• The attraction between molecules is an intermolecular force.
• Intermolecular forces are much weaker than intramolecular forces (e.g. 16 kJ/mol vs. 431 kJ/mol for
HCl).
• When a substance melts or boils the intermolecular forces are broken (not the covalent bonds).
• When a substance condenses intermolecular forces are formed.
Types of The Intermolecular forces
1- Ion-Dipole Forces
• Interaction between an ion (e.g. Na+) and a dipole (e.g. water).
• Strongest of all intermolecular forces:
2- Dipole-Dipole Forces
• Dipole-dipole forces exist between neutral polar molecules.
•
•
Polar molecules need to be close together.
Weaker than ion-dipole forces:
There is a mix of attractive and repulsive dipole-dipole forces as the molecules tumble
If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing
polarity.
3- London Dispersion Forces
• Weakest of all intermolecular forces.
• It is possible for two adjacent neutral molecules to affect each other.
• The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom).
• One instantaneous dipole can induce another instantaneous dipole in an adjacent molecule (or atom).
• The forces between instantaneous dipoles are called London dispersion forces.
• Polarizability is the ease with which an electron cloud can be deformed.
• The larger the molecule (the greater the number of electrons) the more polarizable.
• London dispersion forces increase as molecular weight increases.
• London dispersion forces exist between all molecules.
• London dispersion forces depend on the shape of the molecule.
• London dispersion forces between spherical molecules are lower than between sausage-like molecules.
4- Hydrogen Bonding
As the name "hydrogen bond" implies, one part of the bond involves a hydrogen atom. The hydrogen must be
attached to a strongly electronegative heteroatom, such as oxygen, nitrogen or fluorine, which is called the
hydrogen-bond donor. This electronegative element attracts the electron cloud from around the hydrogen
nucleus and, by decentralizing the cloud, leaves the atom with a positive partial charge. Because of the small
size of hydrogen relative to other atoms and molecules, the resulting charge, though only partial, nevertheless
represents a large charge density. A hydrogen bond results when this strong positive charge density attracts a
lone pair of electrons on another heteroatom, which becomes the hydrogen-bond acceptor.
The most ubiquitous, and perhaps simplest, example of a hydrogen bond is found between water molecules
H-O-H...O-H2
Liquid water's high boiling point is due to the high number of hydrogen bonds each molecule can have
relative to its low molecular mass.
Hydrogen bonding can be found in DNA, RNA, Proteins, alcohols, aldehydes, which have high g higher
boiling points than other compounds
Types of chemical reaction:
Aqueous reactions can be grouped into three general categories
1- Precipitation reactions: in which soluble reactants yield an insoluble solid product, most of such reactions
take
place when the anions and cations of two ionic compounds change partners:
Example:
Pb(NO3)2(aq) + 2KI(aq)  2KNO3(aq) + PbI2(S)
2- Acid-base neutralization reactions: in which an acid reacts with a base to yield water and an ionic
compound called salt,
the driving force here is production of stable water by removal of H+ and OH- ions from the solution.
Example
HCl(aq + NaOH(aq 
H2O(l) + NaCl(aq
3- Oxidation –reduction reactions or redox reactions: in which one or more electrons are transferred
between reaction
Partners (atoms, molecules or ions), the driving force is a decrease in electrical potential.
Group 1A (Li, Na,K,…) and A2 (Ca, Mg, Ba…) elements are strong reducing agents (giving their single
electrons), group
7A elements (F,Cl, Br,I) are strong Oxidizing agent (accepting single electrons)
Example
Mg(s) + I2(g)  MgI2(s)
* Mg gives an electron to each iodine atoms!
Other classification of chemical reactions:
1- Combination reactions: Reactants are fewer than products, this type of reactions most often need burn or
heating to be
performed see below table for some examples.
General Formula A + B  C
One example of combination reaction is major industrial preparation of NH3:
N2(g) + 3H2(g)  2NH3(g)
2- Decomposition reactions: Products are fewer than reactants, this type of reactions energy (most often)
should be
provided in the form of heat to be performed, see below table for some examples.
General Formula C  A + B
3- Displacement Reactions: A reaction in which the atoms or ions of one substance take the place of other
atoms or ion in a
Compound.
General formula
A + BC  AC + B
A number of metals displace hydrogen from acids, such elements are more active than hydrogen
Fe(S) + H2SO4(aq)  H2(g) + FeSO4(aq)
Active metals displace hydrogen from water to produce metal hydroxides
2Li(S) + 2H2O(l)  H2(g) + 2LiOH(aq)
Active metals like Li, K, K, Ba displace weaker metals like CU, Ag, Hg
4- Partner Exchange reactions: In this reaction two products change partners:
General formula
AB + CD  C + BD
Examples:
BaCl2(aq) + H2SO4(aq)  BaSO4(S) +2HCl (aq)
AgNO3(aq) + NaCl(aq) 
AgCl(S) + NaNO3(aq)
Acids and bases
An acid is a substance that provides H+ when dissolves in water.
The most common acids are (HCl (Hydrochloric acid), HNO3 (Nitric Acid), H2SO4 (Sulfuric acid), H3PO4
(Phosphoric acid))
Strong acid is almost completely dissolved in the water: Strong acids are HClO4 (perchloric acid), H2SO4
(Sulfuric acid), HCl (Hydrochloric acid), HNO3 (Nitric Acid)
Week acid is partially dissolved in water: Week acid are: H3PO4 (Phosphoric acid), HF (Hydrofluoric acid),
and CH3CO2H (acetic acid).
Definition of pH
pH is equal to the negative log of the hydrogen ion concentration, or pH = -log [H+]. Using the BrønstedLowry approach that would be pH = -log [H3O+].
The pH of distilled water is 7, this is neutral. Any solution with a pH below 7 (i.e. pH 1.0 to pH 6.9) is an acid
and any solution with a pH above 7 (i.e. pH 7.1 to pH 14) is an alkali.
Acidic solutions have a pH between 1 and 6.9 === your stomach contains HCl it is pH 2.
Alkaline solutions have a pH between 7.1 and 14. === your small intestine is pH 9.
Neutral solutions are neither acidic nor alkaline so their pH is 7.
Some important definitions:
Effective Nuclear charge (Zeff) is the net nuclear charge (attraction) felt by an electron from Nucleus.
- As the average number of screening electrons (electrons closer to nucleus) increases, the effective nuclear
charge decreases.
- As the distance from the nucleus increases, screening electrons increases and Zeff decreases.
Ionization energy (Ei): is the amount of energy which is necessary to remove an electron from an isolated
neutral atom in gaseous state and it is always positive (energy is absorbed). Nobel gases have the highest
ionization energy because electron is held tightly to nucleus, the alkali metals (the first two groups in the
periodic table) have the lowest ionization energy, the above figure show how the Ionization energy increase
through groups and periods Ionization energies INCREASE across a row in the periodic table (left to right)
because of the increasing nuclear charge effect, Ionization energies INCREASE up (bottom to top) a group
because of the decreasing shielding effect
Electron Affinity (Eea): The Energy change that occurs when an electron is added to an isolated atom in
gaseous state and it is most likely negative, because the energy is usually released when a neutral atom adds
an electron.
O(g) + eO-(g) EA1 = -141.3 kJ/mol
O(g) + 2eO2-(g) EA2 = +703 kJ/mol
Group 7 (halogens) have the most negative electron affinity because it have a room in it valence shell for an
additional electron, Group 2 and Nobel gases have electron affinity near zero or positive, the above figure
show how the Ionization energy increase through groups and periods. Electron affinities become greater
(more negative) across a row in the periodic table because of the increasing nuclear charge effect, Electron
affinities become greater (more negative) up a column because of the decreased shielding effect.
Atomic radius: is half the distance between two the nuclei of two identical atoms when they are bonded
together, and it depends on the Effective nuclear charge, the higher the effective Nuclear charge the less the
atomic radius, the above figure show how the atomic radius increase through groups and periods. Atomic
radii increase going down in a group because of the decreased shielding effect, Atomic radii decrease from the
left to right across a row because of the increasing nuclear charge effect.
Electronegativity (EN): The ability of an atom in the molecule (which has a covalent bond or ionic bond) to
attract the shared electrons in the bond example H2O, the two shared electrons between one H atom and O
atom is closer to O atom because it has a higher electronegativity than H, the halogens elements (CL, F…),
Nitogen and oxygen are the most electronegative elements, the alkali metals (the first two groups in the
periodic table are the least electronegative). the above figure doesn’t show how the electronegativity increases
but it is generally similar to ionization energy (increase from left to right and from bottom to top).
electronegativity INCREASE across a row (left to right) because of the increasing nuclear charge effect,
electronegativity INCREASE up (bottom to top)a group because of the decreasing shielding effect
Quantum Mechanical Model of atomic structure and Atomic Orbitals
The Quantum Mechanical models of atomic structure is framed in the form of wave equation (proposed by
the Austrian Physicist Erwin Schrödinger), a mathematical equation similar in form to that used to describe
the motion of ordinary waves in fluid, the solution of wave equation are called wave functions (ψ) or
orbitals.
The wave function ψ gives the shape of the electronic orbital.
The square of the wave function, ψ2 gives the probability of finding the electron in a given volume of space
around the nucleus, that is, gives the electron density for the atom.
A wave function contains three variables called quantum numbers (Principal, Subshell and orbital),
represents as n, l, and ml, which describe the energy level of the orbital and the three dimensional shape and
orientation of the orbital of the region in the space occupied by a given electron.
QUANTUM NUMBERS
1- Principal quantum number(n) n = 1, 2, 3 ….. This is the same as Bohr’s n. As n becomes larger, the
number of the
allowed orbitals are increased and the size of those orbitals becomes larger, the atom becomes larger and
the electron is
further from the nucleus, Principal quantum number also called shell, n = 1 first shell, n = 2 second shell
...etc
2- Subshell quantum number (l), define the three dimensional shape of the orbital, l = 0, 1, 2, ….. (n1),traditionally
was called angular momentum quantum number or azithumal quantum number.
-This quantum number depends on the value of n.
-The values of l begin at 0 and increase to (n - 1). i.e within each shell there are n different number of
shapes for orbitals
- n1 contains one subshell, (l = 0) (“0” also called 1s subshell){one shape s}
- n2 contains two subsells: l = 0, 1 (“0” also called 2s subshell and “1” also called 2p subshell){two shapes
s and p}
- n3 contains three subsells: l = 0, 1, 2 (“0” also called 3s subshell,“1” called 3p subshell, “2” called 3d
subsell) {three
shapes. s, p, and d}
4
n contains four subsells: l = 0, 1, 2, 4 (“0” also called 4s subshell,“1” called 4p subshell, “2” called 4d
subsell, “3” called
4f subsell), {four shapes. s, p, d, and f}
-We usually use letters for l (s, p, d and f )as you see above
-Usually we refer to the s, p, d and f-orbitals.
3- Orbital quantum number(m,) define the spatial orientation of the orbital with respect to a standard set of
coordinate axes
ml = -l,…..0…..+l, so within the same orbital shape there are (2l + 1) different spatial orientation.
Traditionally it was called magnetic quantum number.
This quantum number depends on l, it has integral values between -l and +l. (2l + 1)
when l= 0 then ml = 0, 1s orbital, one spatial orientation,
When l = 1 then ml = -1, 0, +1, 3p orbitals, three spatial orientations,
When l = 2 then ml = -2,-1, 0, +1, +2 (5d orbitals), five spatial orientations,
When l = 3 then ml -3,-2,-1, 0, +1,+2,+3 (7f orbitals), seven spatial orientations ….
4- In addition there is a fourth quantum number (Spin quantum number, (ms)) ms = +1/2 ; -1/2, this means
an electron
can spin either in a clockwise (+1/2)or counterclockwise (-1/2) direction, this discovered by Wolfgang
Pauli in 1926.
so no two electrons in an atom (orbital) can have the same four quantum numbers, if n, l, ml are the same
then ms should be
different (Pauli exclusion principle).
s- orbitals
All s-orbitals are spherical.
As n increases, the s-orbitals get larger.
- As n increases, the number of nodes increase.
- A node is a region in space where the probability of finding an electron is zero. At a node, 2 = 0
- For an s-orbital, the number of nodes is (n - 1).
The p Orbitals
- There are three p-orbitals px, py, and pz.
- The three p-orbitals lie along the x-, y- and z- axes of a Cartesian system. The letters correspond to allowed
values of ml of -1, 0, and +1.
- The orbitals are dumbbell shaped.
- As n increases, the p-orbitals get larger.
- All p-orbitals have a nodal plane at the nucleus.
The d and f Orbitals
There are 5 d-orbitals.
Three of the d-orbitals lie in a plane bisecting the x-, y- and z-axes.
Two of the d-orbitals lie in a plane aligned along the x-, y- and z-axes.
Four of the d-orbitals have four lobes each (cloverleaf shape)
One d-orbital has two lobes and a collar
Note:
Orbitals of the same energy are said to be degenerate.
For n  2, the s- and p-orbitals are no longer degenerate because the electrons interact with each other
Allowed combinations of quantum numbers:
n
l
ml
Orbital
notatio
n
1
0
0
2
0
3
4
Number
of
electrons
in
subshell
2
Number
of orbitals
in shell
Number
of
electrons
in shell
1s
Number
of
orbitals
in
subshell
1
1
2
0
2s
1
2
4
8
1
-1, 0, +1
2p
3
6
0
0
3s
1
2
9
18
1
-1, 0, +1
3p
3
6
2
-2, -1, 0, +1, +2
3d
5
10
0
0
4s
1
2
16
32
1
-1, 0, +1
4p
3
3
2
-2, -1, 0, +1, +2
4d
5
10
3
-3, -2, -1, 0, +1, +2,
+3
4f
7
14
Electron Configurations
Electron configurations tells us in which orbitals the electrons for an element are located
There are three rules to build up electron configurations:
1- Electrons fill orbitals starting with lowest (n) energy orbitals and they move to higher energy orbitals when
the lower energy
orbitals are occupied (lowest energy principle);
2- No two electrons can fill one orbital with the same spin (Pauli’s principle), that is no two electrons can
have the same set
of 4 quantum numbers;
3- For degenerate orbitals (orbitals with the same energy levels like (3p, 5d or7f orbitals), electrons fill each
orbital singly
before any orbital gets a second electron (Hund’s rule).
Electrons filled up in the different orbitals of muliatom as follows:
Periodic table:
Presently there is 114 known elements only 90 of them occur naturally, the remaining have been produced
artificially,
For simplicity each elements has one or two letters, Oxygen O, Nitrogen N, Sodium Na…
Russian chemist Dmitri Mendeleev published the forerunner modern periodic table.
In the modern time periodic table (as you see in the above figure) elements are placed in 7 horizontal rows,
called periods (period 1, 2,3…7) periods also correspond to principle quantum numbers, period 1
correspond to n1, period 2 corresponds to n2……
First period has only two elements (H, He), second and third periods each has 8 elements, fourth and fifth
each has 18 elements, sixth has 32 elements, seventh (incomplete) has 28 elements, the sixth has 32 because
there are 14 elements which are pulled out (Ce, Pr….- LU) (the lower part of the above figure) but there are
belonging to the sixth period, the same thing occur with the seventh period (Th, Pa…L.r) which are belonging
to the seventh period.
Also elements are placed in 18 vertical columns called groups: elements in a given group has similar
chemical properties.
- Group 1 called Group 1A, and group 2 called group 2A both have the s-orbital filled, Group 13- 18 called
Group 3A - 8A
respectively and have the p-orbital filled. Group 1A-8A called main groups.
- Groups 3- 7 are called group 3b-7B respectively, exceptionally group 8, 9, and 10 collectively are called
group 8B, group
11 called group 1B, and group 12 called 2B, all have the d-orbital filled. P.S. the 3d orbital fills after the 4s
orbital.
- All these groups (3B-8B and 1B and 2B ) are called transition metal group.
- The lower two rows elements Collectively called inner transition metal group. (Ce, Pr…...LU) are called
lanthanides
and (Th, Pa…..L.r) are called actinides, both rows have the f- orbital filled. P.S. The 4f orbital fills after the
5d orbital.
Important groups in the periodic table
Group 1A (ns1) also called alkali metal, consists of (Lithium (Li), Sodium (Na), Potassium (K), Rubidium
(Rb)…Fr) all are
shiny, soft metal and react rapidly with water and form alkaline product, hence the name alkali metal,
they never find in
the pure state in the nature, but only in combination with other elements.
Group 2A (ns2) called Alkaline earth metal, consists of (Beryllium (Be), Magnesium (Mg), Calcium
(Ca)….Ra) they are
lustrous, silvery, less reactive than group 1A, never find in the nature in pure state.
Group 3A (ns2, np1) consists of (Boron (B), aluminum (Al), Gallium (Ga), Indium(In), Thallium (Ti)), All are
metals except B
which is semimetal, so all are soft, silvery in appearance, good conductors for heat and electricity, the
most important
and abundant in the earth’s crust is Al.
Group 4A (ns2, np2), the most important elements here are: Carbon (C), (Silicon) (Si) and lead (Pb)
Group 5A (ns2, np3), the most important elements here are: Nitrogen (N), phosphorus (P), and arsenic (As)
Group 6A (ns2, np4), the most important elements here are: Oxygen (O), Sulfur (S)..
Group 7A (ns2, np5) called halogens consists of (Fluorine (F), Chlorine (Cl), Bromine (Br) …At), are
colorful, corrosive
nonmetal, found in nature only in combination with other elements (Group 1A, 2A…)
Group 8A (ns2, np6) called Nobel gases consist of (Helium (He), Neon (Ne), Argon (Ar), Krypton
(Kr)….Rn), are gases
with very low reactivity because the outermost shell is full
Periodic table can be divided into three major classes of elements
1- Metals (most of the elements in the periodic table), all except Mercury (Hg) are solid at the room
temperature good conductor of heat and electricity, Malleable and ductility .
2- Nonmetal, seventeen elements most are located in groups (6A-8A), in addition to Carbon, Nitrogen,
phosphorus and Hydrogen, all are gases t room temperature except Bromine is liquid and Carbon,
Sulfur, Phosphorus and Iodine are solid, they are Brittle, Bright, poor conductors to heat and
electricity.
3- Semimetal, seven elements (Boron (B), Silicon (Si), Arsenic (As), Sb, Ge..), have characteristic in the
middle between
Metals and nonmetals.
Alkáli metals
ns1
Hogens
ns2np5
Noble gases
ns2np6
Change in state of the matter, gaseous, liquid and solid state
Molecular Comparison of Liquids and Solids
Physical properties of substances understood in terms of kinetic molecular theory:
– Gases are highly compressible, assumes shape and volume of container:
• Gas molecules are far apart and do not interact much with each other.
– Liquids are almost incompressible, assume the shape but not the volume of container:
• Liquids molecules are held closer together than gas molecules, but not so rigidly
that the molecules cannot slide past each other.
– Solids are incompressible and have a definite shape and volume:
• Solid molecules are packed closely together. The molecules are so rigidly packed
that they cannot easily slide past each other.
Characteristics of Gases
•
•
No definite shape or volume
vapors
-co-exist with a liquid or solid
•
•
•
•
Gases are highly compressible and occupy the full volume of their containers.
When a gas is subjected to pressure, its volume decreases.
Gases always form homogeneous mixtures with other gases.
Gases only occupy about 0.1 % of the volume of their containers.
Some Properties of Liquids
1- Viscosity
• Viscosity is the resistance of a liquid to flow.
• A liquid flows by sliding molecules over each other.
• The stronger the intermolecular forces, the higher the viscosity.
2- Surface Tension
• Bulk molecules (those in the liquid) are equally attracted to their neighbors.
• Surface molecules are only attracted inwards towards the bulk molecules; therefore surface molecules
are packed more closely than bulk molecules.
• Surface tension is the amount of energy required to increase the surface area of a liquid.
Energy Changes Accompanying Phase Changes
Phase Diagrams
• Phase diagram: plot of pressure vs. Temperature summarizing all equilibria between phases.
• Given a temperature and pressure, phase diagrams tell us which phase will exist.
• Features of a phase diagram:
–
–
–
–
–
Triple point: temperature and pressure at which all three phases are in equilibrium.
Vapor-pressure curve: generally as pressure increases, temperature increases.
Critical point: critical temperature and pressure for the gas.
Melting point curve: as pressure increases, the solid phase is favored if the solid is more dense
than the liquid.
Normal melting point: melting point at 1 atm.
The Phase Diagrams of H2O
–
–
–
–
–
The melting point curve slopes to the left
because ice is less dense than water.
Triple point occurs at 0.0098C and 4.58 mmHg.
Normal melting point is 0C.
Normal boiling point is 100C.
Critical point is 374C and 218 atm.
Solutions and their properties: Concentration of solutions. Ions in aqueous solution:
http://people.wcsu.edu/nigroa/CHE%20111/Lectures/Chapter%2011%20Solutions%20and%20their%20Prop
erties.pdf
Electrolytes and nonelectrolytes
http://www.citycollegiate.com/chapter3b.htm
Chemical equilibrium. The equilibrium constant.
http://library.thinkquest.org/10429/low/equil/equil.htm,
Acids and bases: The pH in solutions of strong acids and strong bases. Equilibria in solutions of weak acids
and weak bases.
http://www.chem.mun.ca/homes/rjhhome/Chem1011/15.1.pdf,
Thermochemistry: Energy changes and energy conservation. Expansion work. Energy and enthalpy.
Hess’s law.
http://s-owl.cengage.com/ebooks/vining_owlbook_prototype/ebook/ch5/Sect5-0.html
Oxidation and reduction: Oxidation state. The activity series of the elements. Balancing redox reactions
http://scienceaid.co.uk/chemistry/physical/redox.html,
DETERMINING OXIDATION STATES
http://pages.towson.edu/ladon/oxstate.html,