Bonding
Forces of attraction that hold atoms together making compounds
Chemical symbols Symbols are used to represent elements Either one capital letter, or a capital letter
with a lower case letter Know names and symbols of elements:
– 1 – 30, plus
–Rb, Cs, Sr, Ba, Ag, Au, Cd, Hg, Pt, Ga, Ge, As, Sn, Pb, Se, Br, I, and U
Basic idea...All chemical bonds form
because they impart stability to the atoms involved
lower energy = greater stability
Quick review All types of chemical bonds
involve electrons Valence electrons, the electrons
in the outermost occupied energy level of an atom, are usually the electrons involved in bonding
The representative elements have the same number of valence electrons as their family number in the American system–Example: Mg, column IIA, 2
valence electrons The transition metals all have
two valence electronsns2(n-1)dx
Lewis dot structures are used to represent the valence electrons–each dot represents a valence
electron
–no more than 8 dots total
–no more than 2 dots on a side
–example = Mg: Na.
.
Lewis dot structures of representative elements
The Octet Rule
Atoms will gain, lose, or share electrons in order to achieve an ns2np6 valence configuration
Sizes of atoms
Periodic trend: atomic radii increase moving down a group– Increasing energy level
Periodic trend: atomic radii decrease moving left to right in a period– The charge felt by the valence electrons
becomes larger
• There is a general decrease in atomic radius from left to right, caused by increasing positive charge in the nucleus.
• Valence electrons are not shielded from the increasing nuclear charge because no additional electrons come between the nucleus and the valence electrons.
Sizes of atoms
• For elements that occur as molecules, the atomic radius is half the distance between nuclei of identical atoms.
• For metals, atomic radius is half the distance between adjacent nuclei in a crystal of the element.
Atomic Radius
Atomic Radius
• Atomic radius generally increases as you move down a group.
• The outermost orbital size increases down a group, making the atom larger.
Sizes of ions
Periodic trend: anions are always larger than the atom they were formed from– Electrons repel each other
Periodic trend: cations are always smaller than the atom they were formed from– Fewer electrons to share same positive
nuclear charge
Ionic Radius
• When atoms lose electrons and form positively charged ions, they always become smaller for two reasons:
1.The loss of a valence electron can leave an empty outer orbital resulting in a small radius.
2.Electrostatic repulsion decreases allowing the electrons to be pulled closer to the radius.
• When atoms gain electrons, they can become larger, because the addition of an electron increases electrostatic repulsion.
Ionic Radius
• Both positive and negative ions increase in size moving down a group.
Ionic Radius
• The ionic radii of positive ions generally decrease from left to right.
• The ionic radii of negative ions generally decrease from left to right, beginning with group 15 or 16.
Ionic Radius
Bonding
Forces of attraction that hold atoms together making compounds
Ionization energy The energy needed to remove a
valence electron from an atom A measure of how tightly the
electrons are being held periodic trend
–increases from the bottom up–increases left to right
In general, metals have lower IE than nonmetals–alkali metals are the lowest IE
family
–noble gases are highest IE family
• The energy required to remove the first electron is called the first ionization energy.
Ionization energy
• First ionization energy increases from left to right across a period.
• First ionization energy decreases down a group because atomic size increases and less energy is required to remove an electron farther from the nucleus.
Ionization energy
Ionization energy
• The ionization at which the large increase in energy occurs is related to the number of valence electrons.
Ionization energy
• Removing the second electron requires more energy, and is called the second ionization energy.
• Each successive ionization requires more energy, but it is not a steady increase.
Ionization energy
Electron affinity A measure of how strongly an
element would like to gain an electron
periodic trend – increases from the bottom up
– increases left to right
– ignore the noble gases
Atoms that lose electrons easily have little attraction for additional electrons (and vice versa)– metals have low IE, low EA– Nonmetals have high IE, high EA
Octet rule: when atoms react, they tend to strive to achieve a configuration having 8 valence electrons
This results in some form of bond formation
Periodic trends… As you move from left to right along a
period… Atoms get
…. Smaller Ionization energy goes
…. Up Electron affinity goes
…. Up
Periodic trends… As you move down a group/family Atoms get
…. Larger Ionization energy goes
…. Down Electron affinity goes
…. Down
Check your understanding
The lowest ionization energy is the ____.
A. first
B. second
C. third
D. fourth
The ionic radius of a negative ion becomes larger when:
A. moving up a group
B. moving right to left across period
C. moving down a group
D. the ion loses electrons
Check your understanding
Electron Configuration of Ions
Na 1s22s22p63s1
– will lose one e- to gain ns2np6 configuration
– Na+ 1s22s22p6
S 1s22s22p63s23p4
– will gain 2 e- to gain ns2np6 configuration– S2- 1s22s22p63s23p6
Ionic Bonding Metals lose electrons easily,
nonmetals have a strong attraction for more electrons
metal atoms will lose electrons to nonmetal atoms, causing both to become ions
1. Metals, having lost one or more electrons, become cations (+)
2. Nonmetals, having gained one or more electrons, become anions (-)
3. Opposites attract: the cations and anions are held together electrostaticly
– called “ionic bonds”
In summary...
Ionic bonds are electrostatic attractions between cations and anions formed when electron(s) are transferred from the low IE, EA metal to the high IE, EA nonmetal
Cation (+)
Ionic compound = crystalline solid
Ionic Compounds
High melting points brittle solids nonconducting as solids conduct electricity as liquids
or aqueous
Ionic Compounds As solids, exist in a 3-D repeating pattern
called a crystal “lattice”
the lattice energy is the energy lowering (stability) accomplished by the formation from “free” ions
Also a measure of the energy required to break apart the ionic compound once formed
The greater the lattice energy, the stronger the force of attraction
Bonding
Forces of attraction that hold atoms together making compounds
Ion dissociation Many ionic compounds will
dissolve in water if it results in more stability (lower E) than in the solid ionic compound
the ions “dissociate” from each other
Ex: CaCl2(s) + H2O Ca2+(aq) + 2Cl-(aq)
Ionic Bond Strength
A measure of the attractive force between the ions
smaller atoms = stronger ionic bonds
fewer atom ratio = stronger bond evidence: melting points
Compare the melting points:
KCl : 776oCKI : 723oCsmaller atoms result in
stronger ionic bonds
Compare the melting points:
CaCl2 : 772oCNaCl : 800oC fewer atoms result in
stronger ionic bonds
Bonding
Forces of attraction that hold atoms together making compounds
Covalent Bonding Covalent bonding involves the
sharing of electron pairs usually between two high EA,
high IE nonmetals–both want more e-’s, neither is
willing to lose the e-’s they have
A nonmetal will form as many covalent bonds as necessary to fulfill the octet rule
example: C, with 4 valence e-’s, will form 4 covalent bonds–results in 8 valence e-’s around
the carbon atom at least part of the time
double and triple covalent bonding is a possibility
When does the octet rule
fail?
H, He and Li Helium strives for 2 valence
electrons– 1s2 configuration
Hydrogen will sometimes will share its one electron with another atom, forming a single covalent bond
Lithium will lose its lone valence electron, gaining the 1s2 configuration of He
Be Be will sometimes lose its 2
valence electrons, gaining the Is2 configuration of He
Be will sometimes form 2 covalent bonds, giving it 4 valence electrons–nuclear charge of +4 cannot handle
8 valence electrons
B Boron will often make three
covalent bonds using its three valence electrons–nuclear charge of +5 cannot
handle 8 valence electrons in a stable manner
“organometallic” compounds
Some metals will form covalent compounds with nonmetals–Hg, Ga, Sn, and others
The octet rule is not followed for the metals,but is for nonmetals
Form 2 or more covalent bonds
P, S, Cl, Se, Br, I Elements in the third period and
lower have empty d orbitals there is room for more than 8
valence electrons These elements will at times
make more than 4 covalent bonds
Rules for Drawing structural formulas
1) Determine the central atom, place the other atoms evenly spaced around the outside
2) Count the total number of valence electrons
3) Draw single bonds between the central atoms and each of the outside atoms
4) Complete the octet on the outside atoms by placing electrons in pairs around the outside atoms (lone pairs)
5) Place any remaining electrons on the central atom in pairs
6) If the central atom does not have its minimum number of electrons (usually 8), form double bonds by moving lone pairs off of the outside atoms and drawing them as bonding pairs
Binary Molecular Nomenclature Two nonmetals no charges to balance multiple subscripts possible
–ex: N2O, NO, NO2, N2O4, N2O5
Use prefixes to represent subscripts mono = 1 di = 2 tri = 3 tetra = 4 penta = 5
Hexa = 6 hepta = 7 octa = 8 nona = 9 deca = 10
Rules, continued..
Change second name to end in “ide”
do not use prefixes on the first word if the prefix is “mono”
always use prefixes on the second name
Examples... CO2
carbon = first word subscript = 1, so no prefix oxide = second word subscript = 2, so prefix = di carbon dioxide
Examples... CO carbon = first word subscript = 1, so no prefix oxide = second word subscript = 1, so prefix = mono carbon monoxide
Examples... SF6
1 sulfur, 6 fluorines sulfur hexafluoride P2O5
2 phosphorus, 5 oxygens diphosphorus pentoxide
Examples... Dinitrogen tetroxide di = 2, so two nitrogen’s tetra = 4, so 4 oxygens N2O4
Dihydrogen monoxide H2O! DHMO.org