Atoms and Bonding-Ch. 2 and 3
Bettelheim text
**Recommended sites: kentchemistry.com
Also, Janet Coonce videos on youtube
• Atom-composed of p+, n0 and e• Know atomic theory!
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Diameter of a nucleus is only about 10-15 m.
Diameter of an atom is only about 10-10 m.
Fig 3.1 The structure of an atom
General, Organic, and Biological Chemistry; Prentice Hall @ 2003,
Ch. 3
Atomic number (Z)
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Always equals the p+ number
In a neutral atom, p+=eMass number: p+ and n0
Isotopes-different number of neutrons, so
different masses
Isotopes of hydrogen
General, Organic, and Biological Chemistry; Prentice Hall @ 2003,
Ch. 3
Periodic Table
• Mendeleev-1869-organized elements by
atomic mass; forerunner of modern PT
• Periods (rows)
• Groups (columns)
• Know names of groups
• Metals, metalloids and non metals
•Fig 3.2 The periodic table
General, Organic, and Biological Chemistry; Prentice Hall @ 2003,
Ch. 3
Periodicity
• Atomic radius-decreases across periods and
increases down groups
• Know characteristics of groups
• Properties of elements determined by earrangement
• Electronic structure (electron configurations)
• *note-shell=principle energy level
• S=spherical orbitals; P=dumbbell shaped
Electronic structure con’t
• The maximum # of e- that can occupy an
energy level (shell) is 2n2
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The first shell has only an s subshell
The second shell has an s and p subshell
The third shell has an s, p, and d subshell.
The fourth shell has an s, p, d, and f
subshell.
General, Organic, and Biological Chemistry; Prentice Hall @ 2003,
Ch. 3
General, Organic, and Biological Chemistry; Prentice Hall @ 2003,
Ch. 3
•Electronic configuration of a few elements are
shown below:
General, Organic, and Biological Chemistry; Prentice Hall @ 2003,
Ch. 3
Sublevel Diagram, Order in which
Sublevels are Usually Filled
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
7p
Electrons in Energy Levels,
Sublevels, and Orbitals
Principal Available
Energy
Sublevels
Level (n) (s, p, d, f)
Orbitals
in Sublevel
(boxes)
Possible
Electrons
in Sublevel
Possible
Electrons in
Energy Level
1
s
1
2
2
2
s
p
1
3
2
6
8
3
s
p
d
1
3
5
2
6
10
18
4
s
p
d
f
1
3
5
7
2
6
10
14
32
Electron configurations con’t
• In class- do configurations for C, N and P
• Also, draw the electron config for
potassium and the potassium ion
• How might an abbreviated configuration of
potassium be written?
General, Organic, and Biological Chemistry; Prentice Hall @ 2003,
Ch. 3
Valence Electrons
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are the electrons in the outermost orbital
of an atom
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are associated with atom’s highest
principal energy level (period)
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only valence electrons determine
chemical properties of an element.
Ionic Bonds
• Typically, atoms from groups IA and IIA
ionically bond with VIA an VIIA
• Ion- Charged atom
The symbol for cation is written by adding a
positive charge as a superscript to the symbol
for the element. For example, by losing an
electron Na metal is converted to the sodium
cation (Na+).
General, Organic, and Biological Chemistry; Prentice Hall @ 2003,
Ch. 4
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Gain of one or more electrons by a neutral atom
gives a negatively charged ion called a anion. The
symbol for the anion is written as by adding a
negative charge as a superscript to the symbol for
the element. For example, by gaining an electron
chlorine atom becomes chloride ion, Cl-.
General, Organic, and Biological Chemistry; Prentice Hall @ 2003,
Ch. 4
•
Halogens: Large ionization energy – electron not easily
lost; Large electron affinity – electron easily gained –
formation of anion is favored.
•
Fig 4.1 Relative ionization energies (red) and electron
affinities (blue) for elements in the first four rows of the
periodic table.
General, Organic, and Biological Chemistry; Prentice Hall @ 2003,
Ch. 4
Where do the electrons come from?
Where do the electrons go?
• Atoms with unfilled valence shells can become
filled by transferring or sharing electrons with
other atoms
• When atoms come together by sharing or
transferring electrons, a chemical bond results.
• After a chemical bond is formed, the atoms
involved in the bond each have a filled valence
shell (octet).
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In NaCl crystal, each sodium ion is
surrounded by 6 chloride ions, and each
chloride ion is surrounded by 6 sodium ions.
In a NaCl solution, individual sodium and
chloride ions are surrounded by water.
+H2O
Notice the
orientation of
the water
molecules
Ionic Bonds
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Opposite electrical charges attract each other.
When sodium combines with chlorine, sodium
transfer electron to chlorine forming Na+ and Clions.
The oppositely charged Na+ and Cl- ions are held
together by a ionic bond.
In three dimensional crystals, all nearby oppositely
charged ions attract each other. We can not separate
attraction between only two oppositely charged
ions-we speak of the whole crystal and crystal as
being an ionic solid or ionic compound. (molecule)
Some Properties of Ionic Compounds
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Ionic compounds are usually crystalline
solids.
Different ions vary in size and charge.
Ions packed together in crystals in different
ways.
The crystal pattern ensures that ions
efficiently fills space and maximizes ionic
bonding.
Ionic compounds have high melting and high
boiling points.
•The formula of an ionic compound shows the
lowest possible ratio of atoms in the compound.
Naming ionic compounds
• We will NOT cover the “old” naming system
• For binary ionic compounds, name the cation first and then
the anion. However, you must change the ending of the
anion to -ide
• If a transition metal is the cation, a roman numeral must be
given to indicate it’s oxidation number (charge)
• Although polyatomic ions bond covalently, they may
ionically bond with a cation or an anion
• For test, know biological importance of Ca, Fe, K and Na
• Polyatomic ions to know for this test: See chart given in
class and know the ones I told you to check [especially
acetate, hydroxide, nitrate, carbonate, sulfate, phosphate,
ammonium, cyanide, chromate, dichromate, hydrogen
carbonate (bicarbonate)]. Quiz Monday!
Practice- Write formulas or
name:
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1. Sodium carbonate
2. Calcium phosphate
3. K2SO4
4. FeCr2O7
Acids and bases
• One definition (Arrhenius):
Acid-substance that donates H+ in water (ex: HCl)
Base- donates OH- (ex: NaOH)
**Bronsted-Lowry definition: Acid is p+ donor; base is p+
acceptor (ex: NH3 + H2O?)
• Hydronium ion H30+
• Exp: HCl, NaOH
• pH scale 0-7-14
• pH + pOH = 14; pH= -log[H+]
• Properties of acids and bases (Acids- react with metals;
“sour”; Bases-slippery, bitter)
• Acid + Hydroxide Base Salt and Water
Acids and bases
• Problem with Arrhenius definition: Not all
acid/ base reactions occur in water
• Bronsted- Lowry definition: Acid is a
proton (H+) donor and base is proton
acceptor
• NH3 + HCl-> NH4+ + Cl-
More acids and bases
• Buffer-substance that can accept hydrogen
or hydroxide ions in solution. Prevent
drastic changes in pH
• Major buffers in the body: 1) bicarbonate
(hydrogen carbonate); 2)hydrogen
phosphate
• CO2 + H2O  H2CO3  H+ + HCO3-
Covalent Bonds
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Covalent bond: A bonds formed by sharing
electrons between atoms. Usually nonmetals
Molecule: A group of atoms held together
by covalent bonds.
The nonmetals near the middle of the
periodic table reach an electron octet by
sharing an appropriate number of electrons.
General, Organic, and Biological Chemistry; Prentice Hall @ 2003,
Ch. 5
•Spherical 1S orbital of two individual hydrogen
atoms bends together and overlap to give an egg
shaped region in the hydrogen molecule. The
shared pair of electrons in a covalent bond is
often represented as a line between atoms.
*Remember, hydrogen doesn’t get a full octet,
it only fills the first shell s orbital
•
A water molecule results when two
hydrogen atoms and one oxygen atom are
covalently bonded in a way shown in the
following picture:
How many electrons are in the
valence shell of each atom in water?
Let’s go back to the electron configurations:
Hydrogen atom: 1 proton, 1 electron.
1s1
Oxygen atom: 8 protons, 8 neutrons, 8 electrons.
1s2 2s2 2p4
Which electronic orbitals are involved in the bonds?
Hydrogen:
1s
Oxygen:
2p
Fig 5.4 For P, S, Cl, and other elements in the third period and
below, the number of covalent bonds may vary, as indicated by
the numbers shown in parentheses.
Numbers of Covalent Bonds
- single bond– sharing of two electrons
- longest bond length
- double bond- sharing of four electrons
- middle bond length
- triple bond- sharing of six electrons
- shortest bond length
- strongest covalent bond
- The electrons in the valence shell that are not
shared are called lone pairs.
Drawing Lewis Structures
•To
draw Lewis structure, you need to know
the connections among atoms.
Also,
common bonding patterns, shown below,
for C, N, O, X, and H simplifies writing Lewis
structure.
Method #1- See next slides and/or handout. I PREFER
THIS WAY
Method #2- Book’s suggested rules for drawing
molecules (Lewis structures)**note-these don’t
always work…they are merely a guide. Some have
expanded octets.
1. Determine the number of valence electrons in the
molecule.
2. Determine the arrangement of atoms in the
molecule (least electronegative –center)
3. Connect the atoms with single bonds
4. Give each peripheral atom an octet by adding lone
pairs
5. Place all remaining electrons on central atom by
adding lone pairs
6. If the central atom still doesn’t have an octet, take a
lone pair from a peripheral atom to form a multiple
bond
LEWIS DOT STRUCTURES
1. Arrange the symbols such that the least
electronegative element is in the center and the other
elements are surrounding the central atom.
O C O
2. Give each of the elements their appropriate number of
valence electrons (dots). Remember the number of
valence electrons for a representative element is the
same as the group number.
..
..
:O
:C: O:
..
..
3. Keep track of the total numbr of valence electrons for
the compound by adding the valence electrons from each
atom. If the compound is an ion then add electrons (dots) for each
negative charge or subtract electrons (dots) for each positive
charge.
4 for C and 6 for O (twice) = 16 electrons
LEWIS DOT STRUCTURES
4. Now move the dots around so that you have 8 dots
(the octet rule) around each element (do not forget the
exceptions) while at the same time keeping the dots in
pairs. Electrons, at this point, exist as pairs (the buddy system).
5. EXCEPTIONS TO THE OCTET RULE: Group I, II, and III
need only 2, 4, and 6 electrons, respectively, around that atom
(if they are covalently bonding).
6. If there are too few pairs to give each atom eight
electrons, change the single bonds between two atoms
to either double or triple bonds by moving the unbonded
pairs of electrons next to a bonding pair.
..
..
:O: :C: : O:
LEWIS DOT STRUCTURES
7. Once the octet rule has been satisfied for each atom
in the molecule then you may replace each pair of dots
between two atoms with a dash.
..
..
:O =C= O:
8. Now check your structure by
a) count the total number of electrons to make sure
you did not lose or gain electrons during the process.
b) Use FORMAL CHARGE (FC) calculations as a
guideline to the correct structure. A zero formal charge
is usually a good indication of a stable structure.
FC (X) = # of valence electrons - (1/2 bonding electrons + nonbonding electrons)
For our example: FC(C) = 4 - (1/2 8 + 0) = 0
FC(O) = 6 - (1/2 4 + 4) = 0
The Basics: Drawing Lewis Structures
Step 1: Calculate the total number of valence electrons in the
molecule or ion
Step 2: Determine the central atom(s) of the molecule or ion –
usually it’s the least electronegative atom.
Step 3: Draw a tentative diagram for the molecule or ion.
Rules
a) A hydrogen atom always forms one bond. Hydrogen
is always a terminal atom in a Lewis diagram – an
atom that is bonded to only one other atom.
b) A carbon atom normally forms four bonds
c) When several carbon atoms appear in the same
molecule, they are often bonded to each other.
Draw the Lewis Structure for the following molecules.
1. H2O
2. CO
Draw the Lewis Structure for the following molecules.
3. BH3
4. NH3
CW- Draw CCl4
CW- Draw PBr3
Draw OCl2
Draw NH4+
*Cation
**Coordinate covalent bond- one atom donates
both shared e-
Explain the last example in terms
of acid/base
• NH3 is the ?
• H+ is the ?
Classwork:
• Do “WS Covalent Lewis Dot Structures”
The next few slides have….
• Some more “challenging” Lewis structures.
I will let you know if we are doing them or
not.
Draw SO3
2-
*anion; name it!
Draw SO2
*Resonant structure. Draw both
Practice:
• Predict the most stable structure:
ONC- or OCN- or NOC-
LEWIS DOT STRUCTURES
Predict the most stable structure: ONC- or OCN- or NOC..
..
..
..
..
..
:O::N::C: or :O::C::N: or :N::O::C:
1) Total electrons is:
6 e- for O + 5 e- for N + 4 e- for C + 1 e- for negative charge = 16 etotal. All structures fulfill the octet rule.
2)
FC (X) = # of valence electrons - (1/2 bonding electrons + nonbonding electrons)
structure#1:
FC(C) = 4 - (1/2 4 + 4) = -2
FC(O) = 6 - (1/2 4 + 4) = 0
FC(N) = 5 -(1/2 8 + 0) = +1
structure#2:
FC(C) = 4 - (1/2 8 + 0) = 0
FC(O) = 6 - (1/2 4 + 4) = 0
FC(N) = 5 -(1/2 4 + 4) = -1
structure #3:
FC(C) = 4 - (1/2 4 + 4) = -2
FC(O) = 6 - (1/2 8 + 0) = +2
FC(N) = 5 -(1/2 4 + 4) = -1
structure #2 has the combination with the lowest
formal charge. It also has the negative formal
charge on one of the more electronegative
atoms. Calculate the formal charge for the most
stable structure:
..
:O:C:::N:
..
(-1, 0, 0)
Practice problems
• Draw: ClO2• SiH4
• AsH3
More Practice Problems
ClO2-
AsH3
SiH4
Group Study Problems
1.
Draw the Lewis structure for the following
a) H2S
b) PH3
c) CH2O
d) NO2-
e) H2CO3
f) CBr4
g) CH2FCl
h) C2H2
I) O3
2. Calculate the formal charge for “c”,
“d”, “f”, and “I”.
So, how do you draw these structures?
* Practice any we didn’t do in class at home
• Br2
• CH4
• HCl
• SiF4
• PF3
• H2CO
• C2H6
• C2H4
• CO2
• CO
• H2O
• OH• NH3
• HCN
• NH4 + (What is a coordinate covalent bond?
• NO3- (*use the ‘rules” for this one. Also,
this one displays resonance. What is that?)
• CO3 2• SO2
Shape of Molecules
•Molecular shapes can be predicted by noting
how many bonds and electron pairs surround
individual atoms and applying what is called the
valence-shell electron-pair repulsion (VSEPR)
model. The basic idea of VSEPR model is that
the negatively charged clouds of electron in
bonds and lone pair repel each other therefore
tends to keep apart as far as possible causing
molecules to assume specific shape.
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There are three step to applying the
VSEPR model:
Step 1: Draw a Lewis structure of the
molecule, and identify the atom whose
geometry is of interest.
Step 2: Count the number of electron charge
clouds surrounding the atom of interest.
Step 3: Predict molecular shape by assuming
that the charge clouds orient in space so that
they are as far away from one another as
possible.
VSEPR Model
The shape depends on the number of charged clouds
surrounding the atom
© 1995-2002 by Prentice-Hall, Inc. A Pearson Company
Bond angles to know:
• Linear-180
• Trigonal planar-120 (If all e- are in bonded pairs
like BF3. If one pair is lone….and thus “bent”, like
SO2 , then it’s actually 119)
• Tetrahedral-109.5
• Pyramidal-107 (lone pair pushes bonded pairs
closer together)
• Bent (water)-104.5 (2 lone pairs)
• Practice:http://www.chem.umass.edu/genchem/wh
elan/Class_Handouts/111_Molecular_Geometry_
Worksheet.pdf
Polar Covalent Bonds and
Electronegativity
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Electrons in a covalent bond occupy the region
between the bonded atoms.
If the atoms are identical, as in H2 and Cl2,
electrons are attracted equally to both atoms
and are shared equally.
If the atoms are not identical, however, as in
HCl, the bonding electrons may be attracted
more strongly by one atom than by the other
and may thus are shared unequally. Such
bonds are known as polar covalent bonds.
• In HCl, electron spend more time near the chlorine
than the hydrogen. Although the molecule is overall
neutral, the chlorine is more negative than the
hydrogen, resulting in partial charges on the atoms.
• Partial charges are represented by placing d- on the
more negative atom and d+ on the more positive
atom.
• Ability of an atom to attract electrons is called the
atom’s electronegativity.
• Fluorine, the most electronegative element, assigned
a value of 4, and less electronegative atoms assigned
lower values
• Electroneg differences >1.7-1.9 considered ionic;
<1.7-1.9 covalent
•Fig 5.7 Electronegativities and the periodic
table
Polar Molecules
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Entire molecule can be polar if electrons
are attracted more strongly to one part of
the molecule than to another.
Molecule’s polarity is due to the sum of
all individual bond polarities (sometimes
called dipolar forces or dipoles) and
lone-pair contribution in the molecule.
•
Molecular polarity is represented by an arrow
pointing at the negative end and is crossed at
the positive end to resemble a positive sign.
•
Molecular polarity depends on the shape of
the molecule as well as the presence of polar
covalent bonds and lone-pairs.
Naming Binary Molecular
Compounds
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When two different elements combines
together they form binary compound.
The formulas of binary compounds are
usually written with the less electronegative
element first. Thus, metals are always
written before non-metals. Prefix such as
mono, di, tri, tetra etc, are used to indicate
number of atoms of each element.
A few examples of binary compounds are
given below:
•
The following two steps guide is
helpful in naming binary compounds:
•
Step 1: Name the first element in the
formula, using a prefix if needed to
indicate the number of atoms.
Step 2: Name the second element in the
formula, using an –ide ending as for
anions, along with a prefix if needed to
indicate the number of atoms.
•
Characteristic of Molecular
Compounds
•Molecules are neutral as a result there is no
strong electrical attractions between the
molecules to hold them together. However,
there are several weaker forces exist between
molecules, known as intermolecular forces
(van der Waals forces).
• When intermolecular forces are very weak,
molecules are weakly attracted to one another
and that the substance is gas at ordinary
temperature.
• If the intermolecular forces are somewhat stronger,
the molecules are pulled together into a liquid.
• If the forces are stronger, the substance becomes
a molecular solid.
• Hydrogen bond-specific type of intermolecular
force
involving hydrogen (that is cov. bonded to a highly
electronegative atom) attracted to an unshared electron pair
of another molecule. Ex-water
• Melting points and boiling points of molecular
solids are lower than those of ionic solids.
• Most molecular compounds are insoluble in
water.
•
Molecular compounds do not conduct electricity when
melted because they have no charged particles.
Remember
• Electronegativity increases across periods
and decreases down groups
• Ionization Energy increases across periods
and decreases down groups
• Atomic radius decreases across periods and
increases down groups
Linear
Bent
Trigonal planar
Tetrahedral