Download Summary of things you need to know from this first quarter. The

Summary of things you need to know
from this first quarter.
The Periodic Table - summary
The modern periodic table contains the elements ordered according to their atomic
number (Z), which corresponds to the number of protons in the nucleus of the atom.
The general form of the periodic table is due to Mendeleev (1869) and is shown below.
Hydrogen is
a nonmetal
The rows of the periodic table are known as periods. Moving from left to right across a
period, elements have an increasing number of electrons in their outer shell (or energy
level). They go from metals on the left to non-metals on the right. The columns of the
periodic table are known as groups. Elements in the same group have the same
number of electrons in their outer shells. Because chemical reactivity depends on the
interaction of electrons in the outer shell, elements in the same group have similar
chemical properties.
An earlier attempt by Newlands (1863) to arrange elements by their relative atomic
masses had a number of drawbacks:
• The positions of some pairs of elements are reversed when ordered by mass (e.g. K
and Ar)
• Not all elements had been discovered at the time and Newlands left no spaces for
undiscovered ones
• Some ‘groups’ contained elements with obviously differing properties (e.g. the group
containing oxygen and sulfur also contained iron!)
Of course, at the time, scientists did not know about protons and electrons (and hence
atomic number) and many elements were, as yet undiscovered.
Group 1 (the alkali metals) – lithium, sodium, potassium, rubidium, caesium
• are soft, reactive metals with low density which must be stored under paraffin
• have atoms with one electron in the outer shell
• react with water to form hydrogen gas and the metal hydroxide (which dissolves in
water to give an alkaline solution),
e.g. potassium + water (hydrogen oxide) _ potassium hydroxide + hydrogen
• become more reactive going down the group as the outer electron (being further from
the nucleus) is less strongly attracted to the positive nucleus and hence more easily lost
Group 7 (the halogens) – fluorine, chlorine, bromine, iodine
• have low melting and boiling points, going from fluorine and chlorine (gases), through
bromine (liquid) to iodine and astatine (solids)
• have atoms with seven electrons in the outer shell
• are all colored, with the colors becoming deeper down the group
• become less reactive going down the group as the outer electrons (being further from
the nucleus) are less strongly attracted to the positive nucleus and hence additional
electrons are less easily gained
• will displace other halogens of lower reactivity from their salts
e.g. chlorine + potassium bromide _ bromine + potassium chloride
Group 0 (the noble gases) – helium, neon, argon, krypton, xenon
• are unreactive gases with full outer shells of electrons
Transition metals These are found in the central area of the periodic table. All other
elements in the periodic table can donate or receive a fixed number of electrons and
therefore make a fixed number of bonds to other elements. Transition metals can
exchange a range of electrons and can therefore change the number of bonds they
make.
Periodic Table Vocabulary
Periodic Table – a chart that organizes information about all of the known elements
according to their properties.
Matter – anything that has mass and volume.
Element – a substance that cannot be broken down into a simpler substance by
ordinary chemical means.
Elemental Symbol – the single capital letter or capital letter followed by a lower case
letter or letters that represents the name of the element.
Atom – the smallest unit of an element that has all of the properties of the element;
basic building block of matter.
Molecule – two or more atoms covalently bond in fixed proportion.
Compound two or more different elements bond in fixed proportion
Atomic Mass – the number of protons and neutrons in the nucleus of one atom of the
element. Located under the elemental symbol (at the bottom of the element box on the
periodic table).
Atomic Number – the number of protons contained in each nucleus of its atoms of the
element. Located over the elemental symbol (at the top of the element box on the
periodic table).
Period – a horizontal row (left to right) in the periodic table.
Group – a vertical column (up and down) on the periodic table.
Reactivity – describes how likely an element is to form bonds with other elements.
Valence Electrons – electrons that are located in the outermost principle energy level
of an atom.
Electron Cloud – the region surrounding the nucleus of an atom where the electrons
are located.
Ion – an atom that has gained or lost one or more electrons.
Ionic Bond – a bond formed by the transfer of electrons (when valence electrons of
one atom are transferred to the outer energy level of another atom). (M with non M)
Covalent Bond – a bond formed when two atoms share valence electrons.
(non M with non M)
The Atomic Theory States:
(1)
All matter is composed of atoms
(2)
Atoms of a given element are identical
(3)
Atoms of different elements combine in simple whole-number ratios to form
chemical compounds.
Metal – an element or substance that conducts heat and electricity.
Nonmetal – an element that does not conduct electricity or heat and is usually a gas at
room temperature.
Metalloid – an element that has some properties of a metal and some properties of a
nonmetal.
Inert – elements and/or compounds that when put together are unable to react
chemically.
The Law of Conservation of Matter – a scientific law that states that during a
chemical reaction, matter cannot be created or destroyed but can be changed into a
different form.
Period law- The chemical properties of elements tends to repeat over specific atomic
number intervals
Law of definite proportions – elements within a compound have specific mass ratios
Isotopes Atoms of the same element but with different numbers of neutrons
1st ionization energy The energy required to remove the most loosely held valence
electron from an atom to form a positive ion.
Electronegativity The tendency of an atom to withdraw electron density from a covalent
bond, between it and another atom.
Exothermic A process that gives out energy
Endothermic A process that requires energy
Periodic Trends in Atomic Properties
A. Ionization Energy - the energy required to remove an electron from an atom
1. Ionization energy increases for successive electrons
2. Ionization energy tends to increase across a period
a.
electrons in the same quantum level do not shield as
effectively as electrons in inner levels
b.
irregularities at half filled and filled sublevels due to extra
repulsion of electrons paired in orbitals, making them easier to remove
3. Helium has the highest First ionization energy
4. Ionization energy decreases with increasing atomic number within a group
because electrons farther from the nucleus are easier to remove
B. Electron Affinity / electronegativity - the energy change associated
with the addition of an electron / the amount an elements pulls on
electrons contained in a covalent bond
1.
Affinity tends to increase across a period
2.
Affinity tends to decrease as you go down in a period
a. electrons farther from the nucleus experience less nuclear
attraction
b. Some irregularities due to repulsive forces in the relatively small
p orbitals
C. Atomic Radius
1 Determination of radius
a. half of the distance between radii in a covalently bonded
covalent atomic radii"
2r
1.
Periodic Trends
Radius decreases across a period because of
increased effective nuclear charge due to decreased shielding
Radius increases down a group because of
addition of principal quantum levels
The trend for increase in Atomic radius is opposite to the other two main
trends IE and electron affinity (and electronegativity)
An Introduction to the Periodic Table
A. Organization
1. Horizontal row is called a "period" (or series)
2. Vertical column is called a "group" or "family"
a. Group 1A - Alkali metals
b. Group 2A - Alkaline earth metals
c. Group 7A - Halogens (Gr, "salt makers")
d. Group 8A - Noble gases
B. Naming Elements 104 and beyond
Nil = 0
un = 1
bi = 2
tri = 3
quad = 4
Pent = 5
hex = 6
sept = 7
oct = 8
enn = 9
Element 109 = un (1) nil(0) enn(9) ium = unnilennium
Notes - Atoms, Molecules and Ions
Fundamental Chemical Laws
A. Law of Conservation of Mass
1. "Mass is neither created nor destroyed"
2. Translation: In ordinary chemical reactions, the total mass of the
reactants is equal to the total mass of the products
B. Law of Definite Proportion
1. "A given compound always contains the same proportions of elements by
mass"
2. Translation: Compounds have an unchanging chemical formula
C. Law of Multiple Proportions
1. "When two elements form a series of compounds, the ratios of the
masses of the second element that combine with one gram of the first
element can always be reduced to small whole numbers
2. Translation: Sometimes two elements can come together in more than one
way, forming compounds with similar, though not identical formulas
Dalton's Atomic Theory
Please read up on Not Taught in class
A. Atomic Theory
1. Each element is made up of tiny particles called atoms
2. The atoms of a given element are identical
3. Chemical compounds are formed when atoms combine with each other. A
given compound always has the same relative numbers and types of atoms
4. Chemical reactions involve reorganizations of the atoms. The atoms
themselves are not changed in a chemical reaction
Early Experiments to Characterize the Atom
A. J.J. Thomson and the Electron
1. Determined the charge to mass ratio of the electron
2. Reasoned that all atoms must contain electrons
3. Reasoned that all atoms must contain positive charges
B. Robert Millikan and the Oil Drop
1. Oil drop experiments determined the charge on an electron
2.
With charge information, and Thomson's charge/mass ratio, he
determined the mass of an electron (9.11 x 10-31 kg)
3 . The Nuclear Atom - Rutherford's Metal Foil Experiment
1. Fired alpha particles at a sheet of thin gold foil
2. Some particles were greatly deflected
a.
Could not have been deflected by electrons or single protons
b.
Must have been deflected by a positively charged object of substantial
mass, this he concluded was the nucleus.
1) Supported concept of a small, central, positive nucleus where
most of the atom's mass was concentrated
2) Disproved Thomson's "plum pudding" model
The Modern View of Atomic Structure: An Introduction
A. Nucleus
1. Protons - positively charged
2. Neutrons - no charge
3. Small size, high density
a. The mass of all of the cars in the United States in an object that
would easily fit in a teaspoon
B. Electrons
1.
Negatively charged
2.
The source of varying reactivity of different elements
3.
Provide most of the atomic volume
C. Atomic Number
1. Number of protons
D. Mass Number
1. Number of protons
+ number of neutrons
Relative atomic mass and Isotopes
Isotopes:. Atoms with the same number of protons (same element) but different numbers
of neutrons
Relative atomic Mass The weighted average mass of atoms of an element
Calculated by
RAM = (Mass1 x % abundance1) + (Mass2 x % abundance2) + (Mass3 x %3)
Abundance: The proportion of
Calculated by
% Abundance2 = (RAM – Mass1) x 100
(Mass2 – Mass1)
Electrons a closer look
Orbital Shapes and Energies
A. Size of orbitals
1. Defined as the surface that contains 90% of the total electron
probability
2. Orbitals of the same shape (s, for instance) grow larger as n increases
B. s Orbitals
1 orbital
max of 2 electrons
1. Spherical shape
2. Nodes (s orbitals of n=2 or greater) a. Internal
regions of zero probability
C. p Orbitals
3 orbitals
max of 6 electrons
1. Two lobes each
2. Occur in levels n=2 and greater
3. Each orbital lies along an axis (2px, 2py, 2pz)
D. d Orbitals
1.
2.
5 orbitals
max of 10 electrons
Occur in levels n=3 and greater
Two fundamental shapes
a. Four orbitals with four lobes each, centered in the plane indicated in the orbital
label
b.
b. Fifth orbital is uniquely shaped - two lobes along the z axis and a belt centered in the xy
plane
dz2
E. f Orbitals
7 orbitals
max of 14 electrons
Occur in levels n=4 and greater Highly complex shapes
Electron configuration
A. Below Complete set of orbitals
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10
Transition metal exceptions
Cr 1s2 2s2 2p6 3s2 3p6 4s1 3d5
Cu 1s2 2s2 2p6 3s2 3p6 4s1 3d10
If it begins with C move one
electron from s to d except for cobalt
because the door is bolted.
d4 and d9 are close to beginning ½
filled or filled so steal an electron
from the s orbital in order to gain
stability
B. Write electron configuration finding the element
Which period is it in ? Which block is it in? how many steps into the block is it?
counting from the left
Try this with Sulfur
answer [Ne] 3s23P4
When writing noble gas abbreviated electron configurations,
don’t forget to include all other blocks that you have to pass
through when counting from the left.
Quantum Numbers
A. Principal Quantum Number (n)
1. Integral values: 1, 2, 3, ....
2. Indicates probable distance from the nucleus
a. Higher numbers = greater distance
b. Greater distance = less tightly bound = higher energy
B. Angular Momentum Quantum (l) Azimuthal
(this was called the "orbital quantum number" in your general chem book)
1. Integral values from 0 to n - 1 for each principal quantum number n
2. Indicates the shape of the atomic orbitals
Table 7.1 Azimuthal quantum numbers and
corresponding atomic orbital numbers
Value of l
Letter used
0
s
1
p
2
d
3
f
4
g
C. Magnetic Quantum Number (ml)
1. Integral values from l to -l, including zero
2. Magnetic quantum number relates to the orientation of the orbital in
space relative to the other orbitals
•
ml: divides the sub shell into individual orbitals which hold the electrons.
Example d subshell is divided into individual orbitals as follows
•
-2
-1
0
+1
+2
D. Spin Quantum Number
1. Up first = + ½ Then down = - ½
Table
n
1
Quantum numbers for the first four levels of orbitals in the
hydrogen atom
l
Orbital
ml
designation
0
1s
0
# of
orbitals
1
2
0
1
2s
2p
0
-1, 0, 1
1
3
3
0
1
2
3s
3p
3d
0
-1, 0, 1
-2, -1, 0, 1, 2
1
3
5
4
0
1
2
3
4s
4p
4d
4f
0
-1, 0, 1
-2, -1, 0, 1, 2
-3, -2, -1, 0, 1, 2, 3
1
3
5
7
Electron Spin and the Pauli Principle
A. Electronic Spin Quantum Number
1. An orbital can hold only two electrons, and they must have opposite
spins
2. Spin can have two values, +1/2 and -1/2
B. Pauli Exclusion Principle (Wolfgang Pauli)
1. "In a given atom no two electrons can have the same set of four
quantum numbers"
The Aufbau Principle and the Periodic Table
A. The Aufbau Principle
1. "As protons are added one by one to the nucleus to build up elements,
electrons are similarly added to these hydrogen-like orbitals
B. Hund's Rule
1. "The lowest energy configuration for an atom is the one having the
maximum number of unpaired electrons allowed by the Pauli principle
in a particular set of degenerate orbitals
Molecules and Ions
A. Chemical Bonding
1. Covalent bonding - Sharing of valence electrons
2. Ionic bonding - Attraction of oppositely charged ions due to a reaction
in which electrons are transferred
3. Metallic bonding –Ions in a sea of electrons
B. Representing Molecules (Covalently bonded)
1. Chemical formula
a. Symbols for atoms and subscripts
1) H20
2) CH4
2. Structural formula
a. Bonds represented by lines
Ball and Stick
Space Filling
C. Ions
1. Ca
+ions
a. Positive ions formed by the loss of electrons
2 . A n i o n s a. Negative ions formed by gaining electrons
D. Ionic Bonding
1. Bond formed by the attraction between oppositely charged ions
2. Ionic bonding forms ionic solids (salts)
3. Ions can be monatomic (one atom) or polyatomic (more than one
atom)
E. Metallic Bonding
1. Occurs between or within metals
2 Ions in a sea of electrons
3 The valence electrons are said to be delocalized (not fixed to any one
atom)
Nomenclature
Naming Simple Compounds
A. Ionic Compounds
1. Positive ion is always named first, negative ion second
Binary ionic compoundsal with non metal
Contain only two elements but may have more than 2 atoms
Always end in ide
Magnesium chloride MgCl2
You were given a list of ions to memorize
Tips for memorizing the polyatomics:
4 Find the "ate" ion (sulfate, for instance)
sulfate = SO4 25 The "ite" ion always has one less oxygen than the "ate" ion sulfite = SO3 2The prefix "per" (think hyper, meaning "above") is used with the "ate" prefix to
indicate one more oxygen than the "ate" ion persulfate = SO5 27 The prefix "hypo" (meaning "under" or "below") is used with the "ite" prefix to
indicate one less oxygen than the "ite" ion hyposulfite = SO22Examples (Just because you can name it doesn't mean it exists!)
6
Perchlorate
ClO4-
Pernitrate
NO4-
Chlorate
ClO3-
Nitrate
NO3-
Chlorite
ClO2-
Nitrite
NO2-
hypochlorite
ClO-
hyponitrite
NO-
2. Metals with more than one oxidation state (transition metals) must have a roman numeral
to indicate the oxidation state
Fe3+ = iron (III)
Mn+2 = manganese (II)
B. Binary Covalent Compounds
1. Must contain two elements, BOTH nonmetals
a. First element
1) full element name
2) prefix only if there is more than one atom
b. Second element
1)
named as if it were an anion (-ide suffix)
2) always gets a prefix
mono - 1
penta - 5
octa - 8
di - 2
hexa - 6
nona - 9
tri - 3
hepta - 7
deca - 10
tetra - 4
Bonding: General Concepts
Types of Chemical Bonds
A. Metallic Bonding Metal with Metal
1. Ions in a sea of electrons
2. Metals combined with metals (called alloys)
3. Electrons are delocalized and free to move in a conduction band
B. Ionic Bonding Metal with nonmetal
1. Electrons are transferred
2. Metals react with nonmetals
3. Ions paired have lower energy (greater stability) than separated ions
C. Covalent Bonds nonmetal with nonmetal
1. Electrons are shared by nuclei
2. Pure covalent (non-polar covalent)
a. Electrons are shared evenly
3. Polar covalent bonds
a. Electrons are shared unequally
b. Atoms end up with fractional charges
(1) δ+ or δElectronegativity
A. Electronegativity
1. The ability of an atom in a molecule to attract shared electrons to itself
B. Electronegativity Trends
1. Electronegativity generally increases across a period (why?)
2. Electronegativity generally decrease within a family (why?)
C. Characterizing bonds
1. Greater electronegativity difference between two elements means less
covalent character and greater ionic character
Any compound that conducts an electric current when melted is an
ionic compound.
Bond Polarity and Dipole Moments
A. Dipolar Molecules
1. Molecules with a somewhat negative end and a somewhat positive end
(a dipole moment)
2. Molecules with preferential orientation in an electric field
+
+
-
-
+
-
3. All diatomic molecules are nonpolar
Eg O2 or H2
B. Molecules with Polar Bonds but no Dipole Moment
1. Linear, radial or tetrahedral symmetry of charge distribution
a. CO 2 - linear
b. CCl 4 – tetrahedral
c. BF 3 – Trigonal planar
Guided practice
Step 1
Identify the central atom
Step 2
Draw a lewis dot diagram
Step 3
Arranging the other atoms around the central atom
Atoms should be attached clockwise in order of mass smallest first (This includes lone pairs)
Step Four find lines of symmetry
In reality this method only
works because drawings
are in 2D. In the 3D real
world polar molecules are
Asymmetric.
Step five if the molecule has more than one line of symmetry then it is not polar
Predicting Formulas of Ionic Compounds
1.
Placement of elements on the periodic table suggests how many
electrons are lost or gained to achieve a noble-gas configuration
a. Group I loses one electron, Group II loses two, Group VI gains two,
Group VII gains one....
2.
We can use the jigsaw pieces to determine formula
a) Each element can be represented as a jigsaw piece
b) Combine jigsaw pieces so that no unused nodes remain
c) Determine how many of each element was used and construct
the formula
Eg
Sodium (Na) with Oxygen (O)
Na
O
Na 2O
Treat polyatomic ions as single jigsaw pieces
Use parenthesis to show how many polyatomic ions are required
Al2(SO4)3 This formula means 2 Aluminum atoms and 3 sulfate polyatomic ions
Cu(OH)2.5H2O means 1 molecule of Cu(OH)2 with 5 waters attached. The
period means that the following number of molecules are attached.
3. Formulas for compounds are balanced so that the total positive ionic
charge is equal to the total negative ionic charge
Total positive = +6
Total negative = -6
C. Sizes of Ions
1.
2.
3.
4.
Anions are larger than the parent atom
Cations are smaller than the parent atom
Ion size increases within a family
Isoelectronic ions
a. Ions with the same number of electrons
b. Size decreases as the nuclear charge Z increases
The trend for the ionic radii is the more negative the bigger, within the same period
Lewis Structures
A. Electrons and Stability
1. "the most important requirement for the formation of a stable
compound is that the atoms achieve noble gas configurations
2. Duet rule
a. Hydrogen, lithium, beryllium, and boron form stable molecules
when they share two electrons (helium configuration)
3. Octet Rule
a. Elements carbon and beyond form stable molecules when they
are surrounded by eight electrons
B. Writing Lewis Structures
1. Rules
a. Add up the TOTAL number of valence electrons from all atoms
b. Use a pair of electrons to form a bond between each pair of
bound atoms. Lines instead of dots are used to indicate each
pair of bonding electrons
c. Arrange the remaining atoms to satisfy the duet rule for
hydrogen and the octet rule for the second row elements
Exceptions to the Octet Rule
A. Boron Trifluoride
1. Note that boron only has six electrons around it
2. BF3 is electron deficient and acts as a Lewis acid (electron pair acceptor)
3. Boron often forms molecules that obey the octet rule
B. Sulfur Hexafluoride
1. Note that sulfur has 12 electrons around it, exeeding the octet rule
2. Sulfur hexafluoride is very stable
3. SF6 fills the 3s and 3p orbitals with 8 of the valence electrons, and places
the other 4 in the higher energy 3d orbital
C. More about the Octet Rule
1. Second row elements C, N, O and F should always obey the octet rule
2. B and Be (second row) often have fewer then eight electrons around
them, and form electron deficient, highly reactive molecules
3. Second row elements never exceed the octet rule
4. Third row and heavier elements often satisfy (or exceed) the octet rule
5. Satisfy the octet rule first. If extra electrons remain, place them on
elements having available d orbitals
Resonance
A. Nitrate ion
1. Experiments show that all N-O bonds are equal
2. A single Lewis structure cannot represent the nitrate ion
3. A resonance structure is drawn by writing the three variant structures,
connected by a double-headed arrow
B. Resonance
1. When more than one valid Lewis structure can be written for
a particular molecule
2. The actual structure is an average of the depicted
resonance structures
3. Must have more than 2 atoms
4. Must have at least one multiple bond and more than two oxygens
and or nitrogens
Finding resonance structures
Step 1 Draw the Lewis diagram for each of the atoms involved
Step 2 Calculate the total number of electrons
Step 3 Write out the molecule with the least electronegative atom in the center
and connect each with one bond
Step 4 For each bond drawn deduct 2 electrons from the total
Step 5 Place the electrons around the outer atoms until full ( don’t forget that
each bond counts as two electrons)
Step 6 Move lone pair electrons from the outer atoms to make bonds with the
central atom until 4 bonds are made
Formal Charge
Formal charge is used to find the most stable structure for a molecule.
The arrangement that has the lowest range of formal charges on its atoms is the
most stable.
FC = Ve – (LP’s + # Bonds)
FC = formal charge
Ve = valence electrons
Hybridization : Combination of orbitals from
different subshells to produce new equal orbitals
Number of electron directions equal number
of combined orbitals
Number of different directions of e- pairs can
be found must match number of orbitals merged
LP = lone pair electrons
The electrons are in 4
different directions so
the nitrogen needed 1 s
combined with 3p’s to
make 4 new sp3
hybridized orbitals
Molecular Structure: The VSEPR Model
A. Valence Shell Electron Pair Repulsion (VSEPR)
1. The structure around a given atom is determined principally by
minimizing electron-pair repulsions
2. Non-bonding and bonding electron pairs will be as far apart as possible
Arrangement of Electron Pairs Around an Atom Yielding Minimum
Repulsion
# of Electron
Shape
Example Lewis dot diagram
Pairs
2
Linear
Bond angle 180o
3
Trigonal Planar
Bond angle 120o
4
Tetrahedral
Bond angle 109.5o
5
Trigonal bipyramidal
Bond angle 120o
Bond angle 180o
6
Octahedral
Bond angle 90o
B. Effect of Unshared Electron Pairs Complete
1. The ideal tetrahedral angle is 109.5°
Comparison of Tetrahedral Bond Angles
Compound
Structure
Angle between
Hydrogens
Methane
4
0 Lone
109.5°
Bonding
pairs
electron
pairs
Ammonia
3
1 lone
bonding pair
pairs
107°
Water
2
2 lone
bonding pairs
pairs
104.5°
2. Lone (unshared) electron pairs require more room than bonding pairs (they
have greater repulsive forces) and tend to compress the angles between
bonding pairs
3. Lone pairs do not cause distortion when bond angles are 120° or greater
4. Substitution of bonding pairs with lone pairs decreases the bond angle
by 2.5o
C. VSEPR and Multiple Bonds
1. For the VSEPR model, multiple bonds count as one effective electron pair
2. When a molecule exhibits resonance, ANY of the resonance structures can
be used to predict the molecular structure using the VSEPR model
D. Molecules Containing No Single Central Atom
1. Apply the principal of distancing shared and unshared electron pairs
2. Look at real 3-dimensional, rotatable models to develop predictive skills