Reducing abilities of group 1 and 2 metals
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I 1 = 419 kJ . I 1 = 590 kJ I 2 = 1145 kJ . Reducing Abilities of Group 1 and 2 Metals. 2 K(s) + 2 H 2 O(l) → 2 K + + 2 OH - + H 2 (g) . Ca(s) + 2 H 2 O(l) → Ca 2+ + 2 OH - + H 2 (g) . Slide 1 of 35. Oxidizing Abilities of the Halogen Elements (Group 17 ). 2 Na + Cl 2 → 2 NaCl.

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Reducing Abilities of Group 1 and 2 Metals

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Reducing abilities of group 1 and 2 metals

I1 = 419 kJ

I1 = 590 kJ

I2 = 1145 kJ

Reducing Abilities of Group 1 and 2 Metals

2 K(s) + 2 H2O(l) → 2 K+ + 2 OH- + H2(g)

Ca(s) + 2 H2O(l) → Ca2+ + 2 OH- + H2(g)

Slide 1 of 35

General Chemistry: Chapter 9


Oxidizing abilities of the halogen elements group 17

Oxidizing Abilities of the Halogen Elements(Group 17)

2 Na + Cl2→ 2 NaCl

Cl2 + 2 I-→ 2 Cl- + I2

Slide 2 of 35

General Chemistry: Chapter 9


Acid base nature of element oxides

Acid-Base Nature of Element Oxides

  • Basic oxides or base anhydrides:

    Li2O(s) + H2O(l) → 2 Li+(aq) + 2 OH-(aq)

  • Acidic oxides or acid anhydrides:

    SO2 (g) + H2O(l) → H2SO3(aq)

  • Na2O and MgO yield basicsolutions

  • Cl2O, SO2 and P4O10 yield acidicsolutions

  • SiO2 dissolves in strong base, acidic oxide.

Slide 3 of 35

General Chemistry: Chapter 9


Class examples

Class Examples:

  • 1. Which of the following atoms and ions are paramagnetic (i.e. have unpaired electrons). Note: An even number of electrons does not indicate that all electrons are paired. (a) He atom, (b) F atom, (c) As atom, (d) F- ion (e) Al3+ ion and (f) Fe atom.

  • 2. Arrange the following in order of increasing atomic radius: (a) Mg, Ba, Be, Sr (b) Rb+, Se2-, Br- and Sr2+ (c] Ca, Rb, F, S (d) Fe, Fe3+, Fe2+.


Class examples cont d

Class Examples – cont’d:

  • 3. Write balanced chemical equations to represent the reactions of the following oxides with water: (a) SO3(g), (b) P4O10(s), (c) BaO(s) and (d) Li2O(s).

  • 4. Arrange the following atoms in order of increasing first ionization energy: (a) Fr, He, K, Br (b) P, As, N, Sb and (c) Sr, F, Si, Cl.

  • 5. Why are transition metal atoms and ions so often paramagnetic?


Class examples1

Class Examples

  • 6. What information does the term “degenerate orbitals” convey?

  • 7. How do a ground state and an excited state electron configuration differ?

  • 8. How many electrons are described using the notation 4p6? How many orbitals does this notation include? What is the shape of the orbitals described using the 4p6 notation?


Class examples2

Class Examples

  • 9. Are all (neutral) atoms having an odd atomic number paramagnetic? Are all atoms having an even atomic number necessarily diamagnetic? Explain.

  • 10. The following electron configurations do not correspond to the ground electronic state of any atom. Why? (a) 1s22s22p64s1, (b) 1s22s22p63s23p63d2 (c) 1s22s22p63s23p63d84s24p2.


Chemical bonding

Chemical Bonding

  • We’ll begin with a few familiar ideas from high school with particular reference to Main Group Elements.

  • Observation:Noble Gases (He, Ne, Ar, Kr..) are stable in monatomic form in all phases (gas, liquid and solid). Other Main Group elements/atoms are not similarly stable in monatomic form. These atoms are held together by covalent and/or ionic bonds.


Chemical reactions bond formations are associated with energy changes

Chemical Reactions/Bond Formations are Associated with Energy Changes

  • Energy is released when two or more isolated atoms form chemical bonds. The simplest examples involve the formation of homonuclear diatomic molecules. For example:

  • H(g) + H(g) → H2(g)ΔH0 = -436 kJ

  • N(g) + N(g) → N2(g)ΔH0 = -946kJ

  • Cl(g) + Cl(g) → Cl2(g) ΔH0 = -242kJ


Chemical reactions bond formations are associated with energy changes1

Chemical Reactions/Bond Formations are Associated with Energy Changes

  • The energy releasedwhen chemical bonds are formed accounts for the stability of chemical compounds. We could reverse the thermochemical equations given on the previous slide and see, for example:

  • H2(g) → H(g) + H(g) ΔH0 = +436 kJ

  • We see that 436kJ of energy must be supplied to break the bonds in one mole of H2 molecules.


Ionic and covalent bonds

Ionic and Covalent Bonds

  • When Main Group elements react, electrons can be transferred (usually from a metal to a nonmetal) to form ionic bonds. In other cases, pairs of electrons can be shared (usually between nonmetal atoms) to form covalent bonds. In both cases valence electrons are somehow “rearranged” when new chemical bonds are formed. Bond “strengths” vary widely from one compound (or element) to the next.


Ionic bonds main group elements

Ionic Bonds – Main Group Elements

  • In the simplest cases ionic bonds are formed when ametal and a nonmetal react.

  • Metal+ Nonmetal → Binary Ionic ↑ ↑

  • Metal atomshave low ionization energies, low electron affinities and low electronegativity.

  • Nonmetalatomshave high ionization energies, high electron affinities and high electronegativity.


Ionic bonds main group elements1

Ionic Bonds – Main Group Elements

  • Metal atom loses one or more electrons to form a positive ion. The positive ion has a noble gas electron configuration.

  • Nonmetal atom gains one or more electrons to form a negative ion. The negative ion usually has a noble gas electron configuration.

  • In the simplest cases we have two monatomic ions. Complications such as polyatomic ions will arise frequently.


Formation of lif from li and f atoms

Formation of LiF from Li and F Atoms

  • Li (1s22s1) → Li+ (1s2) + e-

  • F (1s22s22p5) + e- → F- (1s22s22p6)

  • In this example the Li+ and F- ions have, respectively, the electron configurations of He and Ne.Ionic compounds formed at room temperature are usually solids. Why? The compound formed must be electrically neutral. Why?


Formation of lif from li and f atoms1

Formation of LiF from Li and F Atoms

  • In practice, lithium fluoride would be formed from the reaction of lithium metal and fluorine gas (F2 molecules). The corresponding balanced chemical equation is:

  • 2 Li(s) + F2 (g) → 2 LiF(s)

  • Ionic compounds are solids at anything close to room temperature. The oppositely charged ions are held together (bonded) by very strong electrostatic forces.


Main group binary ionic compounds

Main Group Binary Ionic Compounds


Covalent chemical bonds

Covalent Chemical Bonds

  • The electronic structures of atoms and molecules have many features in common. Individual atoms often possess unpaired electrons. These atoms are usually chemically unstable. Two such atoms can come together to form a molecule with no unpaired electrons. This process can involve the formation of covalent chemical bonds and is highly exothermic.


Covalent chemical bonds1

Covalent Chemical Bonds

  • In isolated atoms most of the electrons will be found in pairs in a number of different atomic orbitals. When atoms combine to form molecules, valence shell electrons are rearranged. Two electrons from different atoms can “pair up” to form a single covalent bond where the bonding molecular orbital is associated with more than one atom.


Covalent and ionic bonds

Covalent and Ionic Bonds

  • Before moving on to Lewis structures lets practice (review) looking at chemical formulas and identifying whether covalent bonds, ionic bonds or bothare important for the compounds under consideration.


Covalent and ionic bonds1

Covalent and Ionic Bonds


Covalent chemical bonds2

Covalent Chemical Bonds

  • We will consider molecular orbitals in more detail shortly. Before doing so we will use Lewis electron dot structures to represent bonding in both covalent and ionic compounds. In these structures valence shell electrons are represented by dots. The rest of the atom – the nucleus and the core (non-valence) electrons are represented using the chemical symbol for the element.


Lewis theory an overview

Lewis Theory: An Overview

  • Valence e- play a fundamental role in chemical bonding.

  • e- transfer leads to ionic bonds.

  • Sharing of e- leads to a covalent bond.

  • e- are transferred or shared to give each atom a noble gas configuration, the octet.

Gilbert Newton Lewis (1875-1946)

General Chemistry: Chapter 10


Lewis symbols and lewis structures

••

••

••

••

••

As

Bi

P

Sb

N

••

••

••

Al

Ar

Se

••

I

••

••

••

••

Lewis Symbols and Lewis Structures

Si

A chemical symbol represents the nucleus and the coree-.

Dots around the symbol represent valence e-.

General Chemistry: Chapter 10


Reducing abilities of group 1 and 2 metals

Lewis structures

General Chemistry: Chapter 10


Covalent bonding an introduction

Covalent Bonding: An Introduction

General Chemistry: Chapter 10


Coordinate covalent bonds

+

H

H

-

••

N

H

Cl

••

••

H

N

H

••

••

H

H

Coordinate Covalent Bonds

Cl

H

FIGURE 10-2

  • Formation of the ammonium ion, NH4+

General Chemistry: Chapter 10


Multiple covalent bonds

Multiple Covalent Bonds

  • In most cases Main Group elements are surrounded by eight valence shell electrons when molecules are formed – octet rule (H is an impt exception!). Such molecules can contain single bonds (one pair of bonding electrons), double bonds (two pairs of bonding electrons and triple bonds as well as lone pairs of electrons.


Lewis structures

Lewis Structures

  • Class Examples: Construct Lewis structures for H2, F2,O2, N2, CO and CN-. In which species are multiple bonds required to satisfy the octet rule? What type(s) of experimental evidence would suggest that the Lewis structures for these species have physical meaning?


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