Chemical Bonding and Structure Notes

Unit 3: Chemical Bonding and Structure

• Focus: What determines the bonding and properties of a substance?

What You Already Know About Chemical Bonds

• There are only 100 different types of atoms, yet millions of substances exist.

• Atoms bond with each other to form new substances.

Part 1: Ionic, Covalent, and Metallic Bonds

Structure 2.1: Ionic Bonding

• Guiding question: What determines the ionic nature and properties of a compound?

• Standard Level (SL) and Higher Level (HL): Approximately 4 hours of study.

Structure 2.1.1: Formation of Ions

• Metal atoms lose electrons to form positive ions called cations.

• Non-metal atoms gain electrons to form negative ions called anions.

• Predicting ion charge from electron configuration.

• Formation of ions with different charges from transition elements.

Structure 3.1: Periodic Table and Ion Charge

• How the position of an element in the periodic table relates to the charge of its ion(s).

• AHL Structure 1.3: Successive ionization energies of transition elements and variable oxidation states.

Structure 2.1.2: Ionic Bond Formation

• Ionic bond: electrostatic attractions between oppositely charged ions.

• Deducing formula and name of ionic compounds from component ions, including polyatomic ions.

• Binary ionic compounds: cation first, then anion with "-ide" suffix.

• Interconverting names and formulas of binary ionic compounds.

• Polyatomic ions to know: ammonium ($NH<em>4^+$), hydroxide ($OH^-$), nitrate ($NO</em>3^-$), hydrogencarbonate ($HCO<em>3^-$), carbonate ($CO</em>3^{2-}$), sulfate ($SO<em>4^{2-}$), phosphate ($PO</em>4^{3-}$$).

Reactivity 3.2: Redox Reactions and Ionic Compounds

• Formation of ionic compounds from elements as a redox reaction.

• AHL Structure 2.2: Formal charge used to predict sulfate structure.

• AHL Reactivity 3.1: Polyatomic anions as conjugate bases of common acids; relationship between stability and conjugate acid's dissociation constant ($K_a$).

Structure 2.1.3: Properties of Ionic Compounds

• Ionic compounds exist as three-dimensional lattice structures represented by empirical formulas.

• Explaining physical properties, including volatility, electrical conductivity, and solubility.

Formation of Sodium Chloride

• Actual reaction: $2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)$

• Changes in properties and bonding during the reaction.

What Are Ions?

Cations:

• Formed by atoms with low ionization energy.

• Elements lose one or more electrons to gain a noble gas electron configuration.

• Examples: $Na^+$, $Mg^{2+}$, $Al^{3+}$, $NH_4^+$, $Fe^{2+}$.

• Metals form positive ions; cats are always positive!

Anions:

• Formed by non-metals that gain electrons to achieve a noble gas configuration.

• Examples: $Cl^-$, $O^{2-}$, $N^{3-}$, $CO_3^{2-}$.

• An(gry) ions are negative!

Ionic Bonds Require the Formation of Ions

• Illustrative atomic and electronic structure diagrams showing sodium and chlorine forming ions.

Transfer of Electrons To Form Ions

• $Na + Cl \rightarrow [Na]^+ + [Cl]^-$

• $1s^22s^22p^63s^1 + 1s^22s^22p^63s^23p^5 \rightarrow 1s^22s^22p^6 + 1s^22s^22p^63s^23p^6$

• Sodium atom transfers an electron to chlorine atom, forming sodium and chloride ions.

What Is An Ionic Bond?

• Involves transfer of one or more electrons from the outer shell of one atom to the outer shell of another.

• Results in formation of positive and negative ions.

• Electrostatic attraction between oppositely-charged ions.

• Ions arranged in a giant lattice in the solid state.

• High lattice enthalpy responsible for many properties of ionic compounds.

NaCl Lattice Models

• Diagrammatic representation of the sodium chloride lattice structure.

Solubility In Water

• $NaCl (s) \rightarrow Na^+ (aq) + Cl^- (aq)$

• NaCl lattice structure dissolving into hydrated sodium and chloride ions.

Physical Properties of Ionic Solids

• What are the physical properties of ionic solids?

Electrical Conductivity

• Requires charged particles which are able to move.

• $NaCl (s) \rightarrow NaCl (aq)$

Ionic Solids Are Brittle

• Explanation of brittleness due to repulsion of like charges when ions are displaced.

Melting Point, Boiling Point, and Volatility

• Volatility: tendency of a substance to vaporize.

• Ionic solids have high boiling points and are thus not volatile.

Summary of Properties of Ionic Compounds

How Ion Charge Corresponds To The Periodic Table

Formulae of Some Monatomic Ions

Polyatomic Ions to Learn

Finding the Formula of A Compound

• Worked example: Aluminum and oxygen.

• Aluminum forms $Al^{3+}$; oxygen forms $O^{2-}$.

• Cross-multiply charges: $Al<em>2O</em>3$

• Balance charges: 2 x $Al^{3+}$ = 6+ and 3 x $O^{2-}$ = 6-

Formula for Ammonium Phosphate

• Ammonium ($NH<em>4^+$) and phosphate ($PO</em>4^{3-}$).

• Balance charges: 3 x $NH<em>4^+$ = 3+ and 1 x $PO</em>4^{3-}$ = 3-

• Formula: $(NH<em>4)</em>3PO_4$

To Do

1. Review the slides and make notes on the key ideas.

2. Attempt the questions on the following slides.

Questions on the Formulae of Ionic Compounds

• Refer to slide 20 for the table in question 1 (Roman numerals show charge).

Answers to Questions on the Formulae of Ionic Compounds

1. lead nitrate,$Pb(NO<em>3)</em>2$
   barium hydroxide, $Ba(OH)<em>2$ potassium hydrogencarbonate, $KHCO</em>3$
   magnesium carbonate, $MgCO<em>3$ copper sulfate, $CuSO</em>4$
   calcium phosphate, $Ca<em>3(PO</em>4)<em>2$ ammonium chloride, $NH</em>4Cl$

2. (a) KBr
   (b) ZnO
   (c) $Na<em>2SO</em>4$
   (d) CuBr
   (e) $Cr<em>2(SO</em>4)<em>3$ (f) $AlH</em>3$

3. (a) tin(II) phosphate
   (b) titanium(IV) sulfate
   (c) manganese(II) hydrogencarbonate
   (d) barium sulfate
   (e) mercury sulfide

4. (a) $Sn^{2+}$
   (b) $Ti^{4+}$
   (c) $Mn^{2+}$
   (d) $Ba^{2+}$
   (e) $Hg^+$

5. $A<em>2B</em>2$

6. Mg 12: electron configuration [Ne]3s2; Br 35: electron configuration [Ar]3d104s24p5. Magnesium loses two electrons, bromine gains one electron each, forming $MgBr_2$ lattice.

Covalent Bonding

Structure 2.2: The Covalent Model

• Guiding question: What determines the covalent nature and properties of a substance?

• Standard Level (SL) and Higher Level (HL): Approximately 10 hours of study.

Structure 2.2.1: Covalent Bond Formation

• Covalent bond: electrostatic attraction between a shared pair of electrons and positively charged nuclei.

• Octet rule: tendency of atoms to gain a valence shell with 8 electrons.

• Deducing Lewis formulas for molecules and ions with up to four electron pairs on each atom.

• Lewis structures show all valence electrons (bonding and non-bonding pairs).

• Electron pairs shown as dots, crosses, or dashes.

• Molecules containing atoms with fewer than an octet of electrons.

• Organic and inorganic examples.

• Nature of science: Limitations of the octet rule.

• Structure 1.3: Why noble gases form covalent bonds less readily.

• Structure 2.1: Why ionic bonds only form between different elements, while covalent bonds can form between atoms of the same element.

Structure 2.2.2: Single, Double, and Triple Bonds

• Single, double, and triple bonds involve one, two, and three shared pairs of electrons respectively.

• Relationship between number of bonds, bond length, and bond strength.

• Reactivity 2.2: Influence of double and triple bonds on reactivity.

Structure 2.2.3: Coordination Bond

• Coordination bond: covalent bond where both electrons of the shared pair come from the same atom.

• Identifying coordination bonds in compounds.

• AHL: Transition element complexes.

• AHL Reactivity 3.4: Lewis acid-base reactions leading to coordination bonds.

What Is A Covalent Bond?

• Covalent bonds are formed when electrons from different atoms are shared so that each atom attains a noble gas configuration, known as the octet rule.

• The bond is due to the electrostatic attraction between shared pairs of electrons and the positively charged nuclei on each side of the electrons.

• Representation of a chlorine molecule.

Lewis Structures

• The Lewis structure shows all the outer (valence shell) electrons in the molecule but does not put the circle around them.

• A pair of electrons can be represented by either a line, a dot and a cross, two dots, or two crosses.

Formation Of Double And Triple Bonds

• Sometimes there are not enough electrons available for all the atoms to achieve an octet.

• When this happens, the atoms will need to share more than one pair of electrons.

• A double bond is formed when two pairs of electrons are shared.

• A triple bond is formed when three pairs of electrons are shared.

How Far Apart Are The Atoms In A Covalent Bond?

• Bond length is the minimum on the potential energy curve, balancing attractive and repulsive forces.

Bond Length

• Bond length is the distance between the two bonded nuclei.

• The distance is larger in $Br<em>2$ than $Cl</em>2$ because bromine atoms are larger.

Bond Strength

• Bond strength is also known as bond enthalpy.

• This is how much energy is needed to break a bond.

• Stronger bonds are shorter.

To Do

• Use the molymod kits to build the structures for these covalent molecules.

• Take a picture of each of your models and then draw their Lewis structures.

Answers

• Lewis structures for various molecules (NH3, CO2, HCN, etc.) and their types of bonds.

To Do

• Use the molymod kits to build the structures for these covalent molecules.

• Make and draw the lewis structure for the following molecules: 1. A methane molecule 2. A nitrogen molecule 3. An ammonium ion 4. A chlorine molecule 5. A carbonate ion

Polar Covalent Bonds

Structure 2.2.5: Bond Polarity

• Bond polarity results from the difference in electronegativities of the bonded atoms.

• Deduce the polar nature of a covalent bond from electronegativity values.

• Bond dipoles can be shown with partial charges or vectors.

• Electronegativity values are given in the data booklet.

• Structure 2.1: Expected properties of ionic compounds in polar covalent compounds.

Structure 2.2.6: Molecular Polarity

• Molecular polarity depends on both bond polarity and molecular geometry.

• Deduce the net dipole moment by considering bond polarity and molecular geometry.

• Examples include species where bond dipoles do and do not cancel each other.

• AHL Structure 3.2: Features of a molecule that make it "infrared (IR) active".

Electronegativity

• Electronegativity: attraction of an atom for a bonding pair of electrons.

• Measured on the Pauling scale (fluorine = 4.0, francium = 0.7).

• Data booklet provides electronegativity values.

Polar Covalent Bonds

• Occur between atoms with electronegativity difference of 0.5 to 1.7.

• Intermediate nature between ionic and covalent.

• Unequal sharing of electrons creates a bond dipole.

• Atoms have partial positive (δ+) and partial negative (δ-) charges.

Ways of Representing Polar Molecules

• δ shows the slightly negative and slightly positive ends of the bond.

• The arrow in dipole notation shows the direction of electron pull.

Electronegativity Difference of the Hydrogen Halides

• H-H would have an electronegativity difference of 0 so would be non-polar.

Models of Molecules

• Using simulations to visualize dipoles in different molecules.

Considering Ionic and Covalent Structures, S2.4.1 part 1.

• Structure 2.4-From models to materials

• Guiding question: What role do bonding and structure have in the design of materials?

• Standard level and higher level: 4 hours

• Structure 2.4.1-Bonding is best described as a continuum between the ionic, covalent and metallic models, and can be represented by a bonding triangle.

• Use bonding models to explain the properties of a material.

Ionic, Covalent or Polar covalent?

• Bond character varies across the periodic table.

• Ionic: Complete transfer of electrons, e.g., $Na^+Cl^-$, ionic lattice.

• Polar covalent: Partial transfer of electrons, unequal sharing of electrons, e.g., HF, HCl, HBr, HI.

• Pure covalent: Equal sharing of electrons, e.g., $Cl_2$, discrete molecules.

Bonding Continuum

• Note this is only a guide and the H–F bond is still considered a covalent bond, despite the large difference in electronegativity (recall that ionic bonding takes place between metallic and non-metallic elements rather than between non-metal elements such as hydrogen and fluorine).

What type of bond?

• Depends on the Periodic Table position!

• Fluorine and francium are in the top right and bottom left corners, respectively.

To do:

1. Make your own notes on the previous slides.

2. Attempt the problems on the next few slides.

Questions on covalent bonds

1. Discuss what happens to the oxygen to oxygen bond length when hydrogen peroxide decomposes to form oxygen gas.
   $2H<em>2O</em>2(aq) \rightarrow 2H<em>2O(l) + O</em>2(g)$

2. Explain why more energy is required to break the carbon to carbon bond in chloroethene, $C<em>2H</em>3Cl$ than to break the carbon to carbon bond in chloroethane, $C<em>2H</em>5Cl$.

3. Aluminum oxide and aluminum iodide are both white solids at room temperature. Explain why the melting point of aluminum oxide (2072 °C) is much higher than the melting point of aluminum iodide (189.4 °C).

Answers to Questions on covalent bonds

1. The bond length decreases. In hydrogen peroxide, H-O-O-H, the 0-0 bond is a single bond, generally longer than the O=0 double bond in oxygen gas. Double bonds are generally shorter than single bonds as the attractive forces increase.

2. Chloroethene contains a carbon to carbon double bond, C=C, which is a stronger bond than the carbon to carbon single bond, C-C in chloroethane. Double bonds are generally stronger than single bonds as the attractive forces increase.

3. The difference in electronegativity between aluminum and oxygen is 1.8 (3.4 - 1.6). The difference in electronegativity between aluminum and iodine is 1.1 (2.7 - 1.6). The bonding in aluminum iodide will be covalent with the Al-I bond being slightly polar whereas the bonding in aluminum oxide is ionic due to the large difference in electronegativity.

How Ionic Are They?

Exercises

7 Which fluoride is the most ionic?
A NaF
B CsF
C MgF₂
D BaF2
8 Which pair of elements reacts most readily?
A Li + $Br<em>2$ B Li + $Cl</em>2$
C K+ $Br<em>2$ D K + $Cl</em>2$
9 You are given two white solids and told that only one of them is an ionic compound. Describe three tests you could carry out to determine which it is.

Answers To How Ionic Are They

Questions

7 B 8 D
9 Test the melting point: ionic solids have high melting points.
Test the solubility: ionic compounds usually dissolve in water but not in hexane.
Test the conductivity: ionic compounds in aqueous solution are good conductors.

Exercises

Ionic/ covalent questions

10 Which substance contains only ionic bonds?
A $NaNO<em>3$ B $H</em>3COCH<em>3$ C $NH</em>4Cl$
D $CaCl<em>2$ 11 Which of the following molecules contains the shortest bond between carbon and oxygen? A $CO</em>2$
B $H<em>3PO</em>4$
C CO
D $CH<em>3COOH$ 12 For each of these molecules, identify any polar bonds and label them using δ+ and δ- appropriately: (a) HBr (b) $CO</em>2$
(c) ClF
(d) $O<em>2$ (e) $NH</em>3$
13 Use the electronegativity values in Section 8 of the IB data booklet to predict which bond in each of the following pairs is more polar.
(a) C-H or C-CI
(b) Si-Li or Si-Cl
(c) N-Cl or N-Mg

Ionic/ Covalent Answers

10 D 11 C
12 (a)
sigma+H-Brsigma-
(b) sigma-O=C=Osigma+
(c) sigma+Cl-Fsigma-
(d) O=O
(e) sigma+H-Nsigma--H
H
13 (a) C 2.6 H 2.2 difference = 0.4
C 2.6 CI 3.2 difference = 0.6, more polar
(b) Si 1.9 Li 1.0 difference = 0.9
Si 1.9 CI 3.2 difference = 1.3, more polar
(c) N 3.0 CI 3.2 difference = 0.2
N 3.0 Mg 1.3 difference = 1.7, more polar

Structure 2.3: The Metallic Model

• Guiding question: What determines the metallic nature and properties of an element?

• Standard level and higher level: 2 hours

Structure 2.3.1: Metallic Bond

• A metallic bond is the electrostatic attraction between a lattice of cations and delocalized electrons.

• Explain the electrical conductivity, thermal conductivity, and malleability of metals.

• Relate characteristic properties of metals to their uses.

• Tool 1, Inquiry 2, Structure 3.1: Experimental data demonstrating physical properties of metals and trends in the periodic table.

• Reactivity 3.2: Trends in reactivity of metals predicted from the periodic table.

Structure 2.3.2: Strength of Metallic Bond

• The strength of a metallic bond depends on the charge of the ions and the radius of the metal ion.

• Explain trends in melting points of s and p block metals.

• A simple treatment in terms of charge of cations and electron density is required. Structure 2.4: What are the features of metallic bonding that make it possible for metals to form alloys?

METALLIC BONDING

Metallic Bonding

• Involves a lattice of positive ions surrounded by delocalized electrons

• Metal atoms achieve stability by “off-loading” electrons to attain the electronic structure of the nearest noble gas.

• These electrons join up to form a mobile cloud which prevents the newly-formed positive ions from flying apart due to repulsion between similar charges.

Metallic Bonding

• Atoms arrange in regular close packed 3-dimensional crystal lattices.

Metallic Bonding

• The outer shell electrons of each atom leave to join a mobile “cloud” or “sea” of electrons which can roam throughout the metal.

• The electron cloud binds the newly-formed positive ions together.

METALLIC BOND STRENGTH

• Depends on the number of outer electrons donated to the cloud and the size of the metal atom/ion.

• The strength of the metallic bonding in sodium is relatively weak because each atom donates one electron to the cloud. Na

Metallic Bond Strength

• The metallic bonding in potassium is weaker than in sodium because the resulting ion is larger and the electron cloud has a bigger volume to cover so is less effective at holding the ions together. Na K

Metallic Bond Strength

• Depends on the number of outer electrons donated to the cloud and the size of the metal atom/ion.

• The metallic bonding in magnesium is stronger than in sodium because each atom has donated two electrons to the cloud. The greater the electron density holds the ions together more strongly. Na Mg K

METALLIC PROPERTIES

• MOBILE ELECTRON CLOUD ALLOWS THE CONDUCTION OF ELECTRICITY For a substance to conduct electricity it must have mobile ions or electrons. Because the ELECTRON CLOUD IS MOBILE, electrons are free to move throughout its structure. Electrons attracted to the positive end are replaced by those entering from the negative end. Metals are excellent conductors of electricity

Delocalized Electrons

• Metals tend to have a low number of outer (valence) electrons.

• Although the process of removing these electrons is endothermic, generally the energy required is relatively low compared to non-metals.

• Once the electrons become detached from the atoms, they are delocalized (free to move). Metallic bonding is sometimes described as a close-packed lattice of metal cations in a "sea" of delocalised electrons.

Electrical Conductivity

• Electrical conductivity is due to the movement of charged particles. Metals are good conductors of electricity as the current is carried by the delocalised electrons which can move freely through the metal structure (also true for the non-metal graphite) so metals have very low electrical resistance. For the same reason metals are also good conductors of heat.

Melting points of metals

• Most metals have quite high melting points which suggests that metallic bonding is generally quite strong.

• In fact the melting points of metals range from - 38.8 °C to + 3422 °C.

• Mercury (melting point: - 38.8 °C) is the only metal that is a liquid at room temperature (although caesium melts at 28.4 °C.). The metal with the highest melting point is tungsten which melts at 3422 °C.

Malleability and ductility

• Malleable can be beaten into shape with a hammer (the opposite is brittle).

• Ductile can be drawn into a wire.

• Metals are malleable and ductile as the close-packed layers of cations can slide over each other without breaking more bonds than are formed.

Trends in melting points of metals

• When a series of metals forms cations with the same charge then the larger the cation the weaker the metallic bond. As the charge density of the cations decreases the attraction between the cations and the delocalized electrons also decreases. This is clearly shown in the trend in melting points for the group 1 metals.
  *
  Li
  Na
  K
  Rb
  Cs
  M. Pt. / °C 180.5
  97.8
  63.4 39.3
  28.4

Alloys

• Structure 2.4.3-Alloys are mixtures of a metal and other metals or non-metals. They have enhanced properties.

• Explain the properties of alloys in terms of non-directional bonding.

• Illustrate with common examples such as bronze, brass, and stainless steel. Structure 1.1-Why are alloys more correctly described as mixtures rather than as compounds?

Alloys

Some examples of alloys:
Brass (Cu & Zn)
Steel (Fe & C & other metals)
Bronze (Cu & Sn)