1. No category

Chemistry 11 – Unit 5 Forero Name: ______________________________ Date: ___________ Blk: _____
NOTES: BONDING
Examine your periodic table to answer these questions and fill-in-the-blanks. Use
drawings to support your answers where needed:
I. IONIC BONDING
Ionic bond: formed by the attraction of positive ions to negative ions.
1. Examine the Lewis structure examples above. When is an ionic bond formed?
An ionic bond is formed when an electron from one atom is transferred to another
atom, so as to create one positive ion and one negative ion.
2. All the electrostatic attraction in an ionic bond depends on the distance separating the
central charges: the greater the distance, the smaller the attraction.
3. What TWO trends found in the periodic table are responsible for ionic bonds?
Electronegativity and ionization energy: atoms with high electronegativity and
ionization energy form ionic bonds with atoms with low electronegativity and
ionization energy.
4. How to predict when an ionic bond will form:
IONIC BONDS are formed when elements from opposite sides of the
periodic table are combined à when a metal and non-metal are combined.
5. Which of the following atom pairs would you expect to form ionic bonds?
(a) Ba and S (b) P and Cl
(c) Ca and O (d) Rb and I
(e) O and H
(f) S and O
Chemistry 11 – Unit 5 Forero The strength of ionic bonds can be estimated by examining the MELTING
TEMPERATURES of ionic compounds.
6. Look at the melting temperatures of some ionic substances:
- LiF = 845°C
- KCl = 770°C
- LiCl = 605°C
- NaF = 993°C
What can you conclude about the strength of ionic bonds? Why?
Ionic bonds are very strong because it takes a great deal of energy to break them.
Recall what you know about the trend of atomic radius on the periodic table to answer the
following questions.
7. (a) Which compound has the smaller distance between the nuclei of the two ions
involved: NaCl or KBr?
NaCl
(b) What happens to the force of electrostatic attraction between the two ions in an ionic
bond as the ions get smaller?
The smaller the ions, the smaller the distance between the + and – charges and the
greater the force of attraction between the ions.
(c) What happens to the strength of an ionic bond as the ions involved gets smaller?
What happens to the melting temperature?
The smaller the ions involved, the greater the ionic bond strength and the higher
the melting temperature.
8. Mg2+ and Na+ have roughly the same ionic radius. O2- and F- have roughly the same
ionic radius. Which substance should have a higher melting temperature: NaF or MgO?
Although the ions are about the same size, there is more charge on both O2- and
Mg2+. Recall that the greater the charge, the greater the force of attraction. The
increased force of attraction requires greater energy to separate the ions and
therefore a higher melting temperature.
9. Which member of each of the following pairs would you expect to have the higher
melting temperature?
i. CaO or RbI
ii. BeO or BN
iii. LiF or NaCl
iv. CsCl or BaS
v. RbI or KCl
vi. BeO or Mg
Chemistry 11 – Unit 5 Forero II. ION SIZE
Negative ions: assume extra electrons are added to a neutral atom of O to make O2-.
The resulting ion has the same positive nuclear charge and an increased number of
negative electrons surrounding the nucleus (recall the fire/shading analogy).
1. What happens to the amount of electrostatic repulsion existing between the electrons?
The amount of repulsion increases.
2. What happens to the volume occupied by the electrons due to the change in the amount
of electron-electron repulsion?
The volume increases.
3. Fill in the appropriate word:
NEGATIVE IONS are LARGER than the corresponding neutral atom.
Positive ions: assume extra electrons are removed to a neutral atom of Mg to make Mg2+.
The resulting ion has the same positive nuclear charge and a decreased number of
negative electrons surrounding the nucleus.
1. What happens to the amount of electrostatic repulsion existing between the electrons?
The amount of repulsion decreases.
2. What happens to the volume occupied by the electrons due to the change in the amount
of electron-electron repulsion?
The volume decreases.
3. Fill in the appropriate word:
NEGATIVE IONS are SMALLER than the corresponding neutral atom.
4. Examine the diagram below, which shows a section of crystal NaCl.
Which circles represent Na+: the larger or smaller ones?
Chemistry 11 – Unit 5 Forero III. COVALENT BONDING
Covalent bond: A bond that involves the equal sharing of electrons.
Octet Rule: Atoms in families 14 to 17 of the periodic table tend to form covalent bonds
so as to have eight electrons in their valence shells.
1. What trend found in the periodic table is responsible for ionic bonds?
Electronegativity: covalent bonds form when both atoms involved have relatively
large electronegativities, attract each other’s electron strongly and will not let go
of their own electrons.
2. How to predict when an ionic bond will form:
COVALENT BONDS are formed when a non-metal combines with another
non-metal.
3. Which of the following atom pairs would you expect to form ionic bonds?
(a) S and O
(b) Ba and O (c) Fe and Cl (d) N and O
(e) H and S
(f) C and H
4. Look at the melting temperatures of some covalent crystals:
- BN = 3000°C
- SiC = 2700°C
- C (diamond) = 3550°C
What can you conclude about the strength of covalent bonds? Why?
Covalent bonds are very strong because it takes a great deal of energy to break
them.
It is very tempting to say that “covalent compounds have high melting temperatures”, but
look at the melting temperatures of the following covalent compounds:
- CH4 = -182°C
- O2 = -218°C
- HCl = -114°C
Chemistry 11 – Unit 5 Forero 5. Why do you think these covalent compounds have such low melting points?
The crystals in question 6 are being held together by a network of covalent
bonds extending between every atom in the crystal. This network makes each
crystal one huge “molecule” held together by identical covalent bonds, yielding a
high melting temperature.
The compounds in this question consist of individual molecules, which contain
covalent bonds. The individual molecules of CH4, O2 and HCl in a solid are held
next to each other by much weaker bonds.
IV. LONDON FORCES
Individual molecules are held together by strong covalent bonds between the atoms in the
molecule. Such bonds are called INTRAMOLECULAR FORCES (intra = within).
There are weak forces that hold one complete, neutral molecule next to another such
molecule. These INTERMOLECULAR FORCES (inter = between) are called LONDON
FORCES.
The London dispersion force is the weakest intermolecular force.
Dipole: partial separation of charge which exists when one end of a molecule has a slight
excess of positive charge and the other end of the molecule has a slight excess of negative
charge.
London force: a temporary attractive force that results when the electrons in two adjacent
atoms occupy positions that make the atoms form temporary dipoles.
London forces are the attractive forces that cause nonpolar substances to condense to
liquids and to freeze into solids when the temperature is lowered sufficiently.
Chemistry 11 – Unit 5 Forero Because of the constant motion of the electrons, an atom or molecule can develop a
temporary (instantaneous) dipole when its electrons are distributed unsymmetrically
about the nucleus.
A second atom or molecule, in turn, can be distorted by the appearance of the dipole in
the first atom or molecule (because electrons repel one another), which leads to an
electrostatic attraction between the two atoms or molecules.
Dispersion forces are present between all molecules, whether they are polar or nonpolar.
The more electrons an atom or molecule has altogether, the stronger the London forces
existing between I and a neighbouring atom or molecule.
1. What happens to the strength of the London forces between two identical atoms going:
(a) down a family of the periodic table? Why?
Increases. Larger and heavier atoms and molecules exhibit stronger dispersion
forces than smaller and lighter ones.
(b) left to right across a period of the periodic table? Why?
Decreases. In a smaller atom or molecule, the valence electrons are, on average,
closer from the nuclei than in a smaller atom or molecule. They are more tightly
held and can less easily form temporary dipoles.
The ease with which the electron distribution around an atom or molecule can be
distorted is called the polarizability.
2. In response to a request to “write an equation showing what happens when H2(s) melts”,
a student writes the following:
H2(s) à 2 H(l)
What does this equation incorrectly imply about the bonds & forces in a sample of H2(s)?