MYP 10 Chemistry 2013-14
Bonding Worksheet
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Class: _________ Date: _____________
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4.1 Ionic bonding
4.1.1 Describe the ionic bond as the electrostatic attraction between oppositely charged ions.
4.1.2 Describe how ions can be formed as a result of electron transfer.
4.1.3 Deduce which ions will be formed when elements in groups 5,6 and 7 gain electrons.
4.1.4 Deduce which ions will be formed when elements in groups 5,6 and 7 gain electrons.
4.1.5 State that transition elements can form more than one ion.
4.2 Covalent bonding
4.2.1 Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei.
4.2.2 Describe how the covalent bond is formed as a result of electron sharing.
4.2.3 Deduce the Lewis (electron dot) structures of molecules and ions for up to four electron pairs on each atom.
4.2.4.State and explain the relationship between the number of bonds, bond length and bond strength.
4.2.5 Predict whether a compound of two elements would be covalent from the position of the elements in the periodic table or
from their electronegativity values.
4.2.6 Predict the relative polarity of bonds from electronegativity values.
4.2.7 Predict the shape and bond angles for species with four, three and two negative charge centres on the central atom using the
valence shell electron pair repulsion theory (VSEPR).
4.2.8 Predict whether or not a molecule is polar from its molecular shape and bond polarities.
4.2.9 Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite and C60 fullerene).
4.2.10 Describe the structure of and bonding in silicon and silicon dioxide.
4.3 Intermolecular forces
4.3.1 Describe the types of intermolecular forces (attractions between molecules that have temporary diploes, permanent dipoles
or hydrogen bonding) and explain how they arise from the structural features of molecules.
4.3.2 Describe and explain how intermolecular forces affect the boiling points of substances.
4.4 Metallic bonding
4.4.1 Describe the metallic bond as the electrostatic attraction between a lattice of positive ions and delocalized electrons.
4.4.2 Explain the electrical conductivity and malleability of metals.
4.5 Physical properties
4.5.1 Compare and explain the properties of substances resulting from different types of bonding.
1. Write the formula for each of the following compounds:
(a) sodium nitrate
(b) lead(II) nitrate
(c) ammonium phosphate
(e) aluminium hydride
2. Name the following compounds: (a) Sn3(PO4)2
3. Which fluoride is most ionic?
A. NaF
(b) Mn(HCO3)2
B. CsF
C. MgF2
(d) chromium(III)sulphate
(c) Hg2S
(d) Ti(SO4)2
D. BaF2
4. The letters A, B, C and D represent four consecutive elements in the periodic table. The number of
electrons in the highest occupied energy levels are
A: 3
B: 4
C: 5
D: 6
Write the formula for
(a) An ionic compound that form between A and C, showing the charges.
(b) A covalent compound containing B and D.
5. Draw Lewis structures for HCl, Cl2, ClF3 and ClF5. In which of the molecules does chlorine obey octet
rule?
6.
It is somewhat artificial to classify bonds based on the differences in the electronegativities (X) of the two
atoms. However, we will use these ranges to do so:
Ionic
X > 1.7
(symbolized as A+ and Z-)
Polar Covalent
1.7  X
(symbolized as A+ and Z-)
Non-polar Covalent
X = 0
(no charges)
(a) For each of these molecules, identify any polar bonds and label the atoms using + and - appropriately.
(i) HBr
(ii) CO2
(iii) ClF
(iv) CuO
(v) O2
(vi) NH3
(b) Use the electronegativity values in Table 7 of the Data Booklet, predict which bond in each of the
following pairs is more polar:
(i) C – H or C – Cl
(ii) Si – Li or Si – Cl
(iii) N – Cl or N – Mg
(c) Use the electronegativity values in Table 7 of the Data Booklet, determine the type of bonding in each of
the following compounds. Where possible, state whether the molecule is polar or non-polar.
(i) AlF3
(ii) P4
(iii) CS2
(iv) C2H6
7. Predict the shape and bond angles of the following molecules and ions:
(a) H2S
(b) CF4
(c) HCN
(d) NF3
(e) CO32-
8. How many negative charge centres are there around the central atom in molecules that have the
following shapes?
(a) tetrahedral
(b) bent
(c) linear
(d) trigonal pyramidal
(e) trigonal planar
9. Predict the shape of each of the following ions:
(a) NO3-
(b) NO2+
(c) NO2-
10(a) Use the VSEPR theory to deduce the shape of H3O+ and C2H4. For each species, draw the Lewis
structure, name the shape and state the value of the bond angle(s).
(b) Predict and explain whether each species is polar.
(c) Using Table 7 of the Data Booklet, predict and explain which of the bonds O-H, O-N or N-H would
be most polar.
11. Describe the shapes of the four fluorides, BF3 , CF4, NF3 and SF6.
Explain their shapes in terms of the number and types of electron pairs they contain [VSEPR theory].
12. Explain at the molecular level why ethanol (C2H5OH) is soluble in water, but cholesterol (C27H45OH)
and ethane (C2H6) are not.
[SL paper 2, May 01]
13. For each of the following compounds PCl3, PCl5 and POCl3
(a) Draw a Lewis structure for each molecule in the gas phase.
(b) State the shape of each molecule and predict the bond angle.
(c) Deduce whether or not each molecule is polar, giving a reason for your answer.
14. Explain why carbon dioxide is non-polar molecule but sulphur dioxide is polar.
15. Explain the following in terms of structure and bonding.
(a) Sodium oxide has a high melting point, does not conduct electricity when solid but conducts electricity
when molten.
(b) Phosphine, PH3, has a lower boiling point than ammonia, NH3, and arsine, AsH3.