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 ions to ions. 1. Examine the Lewis structure examples above. When is an ionic bond formed? 2. All the electrostatic attraction in an ionic bond depends on the separating the central charges: the greater the, the the attraction. 3. What TWO trends found in the periodic table are responsible for ionic bonds? 4. How to predict when an ionic bond will form: IONIC BONDS are formed when elements from of the periodic table are combined à when a and 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 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? 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? (b) What happens to the force of electrostatic attraction between the two ions in an ionic bond as the ions get smaller? (c) What happens to the strength of an ionic bond as the ions involved gets smaller? What happens to the melting temperature? 8. Mg 2+ and Na + have roughly the same ionic radius. O 2- and F - have roughly the same ionic radius. Which substance should have a higher melting temperature: NaF or MgO?
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 II. ION SIZE iii. LiF or NaCl iv. CsCl or BaS v. RbI or KCl vi. BeO or Mg Negative ions: assume extra electrons are added to a neutral atom of O to make O 2-. 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? 2. What happens to the volume occupied by the electrons due to the change in the amount of electron-electron repulsion? 3. Fill in the appropriate word: NEGATIVE IONS are than the corresponding neutral atom. Positive ions: assume extra electrons are removed to a neutral atom of Mg to make Mg 2+. 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? 2. What happens to the volume occupied by the electrons due to the change in the amount of electron-electron repulsion? 3. Fill in the appropriate word: NEGATIVE IONS are 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?
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? 2. How to predict when an ionic bond will form: COVALENT BONDS are formed when a combines with another. 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? It is very tempting to say that covalent compounds have high melting temperatures, but look at the melting temperatures of the following covalent compounds: - CH 4 = -182 C - O 2 = -218 C - HCl = -114 C
5. Why do you think these covalent compounds have such low melting points? 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. 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? (b) left to right across a period of the periodic table? Why? 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 H 2(s) melts, a student writes the following: H 2(s) à 2 H (l) What does this equation incorrectly imply about the bonds & forces in a sample of H 2(s)?