PARTICLE INTERACTIONS IN CHEMISTRY – UNIT REVIEW

Prior Knowledge

• Balancing chemical equations (Y9)

• Chemical Nomenclature (Y9)

• Draw Lewis dot diagram for the first 20 elements and all main group elements (Y9)

• Distinguish between ionic and molecular compounds in terms of metals and non-metals (Y9)

• Distinguish between chemical and physical properties, chemical and physical changes (Y8)

• Define atom, ion, valence electron, polyatomic ion (Y8)

New Knowledge

• Define and describe electronegativity.

• Describe the periodic trends of atomic radius and electronegativity, giving reason for the similarities and differences.

• Calculate electronegativity difference (ΔEN) and explain its relevance to chemical structure.

• Distinguish between ionic and covalent bonds in terms of electron distribution.

• Distinguish between polar and nonpolar covalent bonds in terms of electron distribution.

• Define and describe a dipole.

• Explain how a bond can result in a dipole.

• Draw complete Lewis structures (including complete and partial charges/dipoles where relevant) for ionic formula units, molecules, and polyatomic ions.

• Distinguish between bond polarity and molecular polarity, including describing the 3-dimensional symmetry of a molecule.

• Describe the physical and chemical properties of ionic and molecular compounds.

• Explain how the properties of a substance can be determined by intra- or intermolecular interactions.

• Describe each of the intermolecular forces studied in this course:

  • London Dispersion Forces

  • Dipole-Dipole

  • Hydrogen Bonding

  • Ion-Dipole

• Explain the factors that affect physical properties of matter (e.g. state of matter, melting point, boiling point, density, solubility, etc.) including intermolecular forces, molecule size, and temperature.

• Describe the principle of minimum potential energy and its relevance to the formation of bonds.

• Explain the role kinetic (thermal) energy plays in the physical properties of a substance.

• Interpret data based on knowledge of attractive forces and potential energy between ions in a lattice or molecules in a substance.

Energy in Chemistry

There is a fundamental principle in physics and chemistry which states that systems will tend towards their lowest state of potential energy. We have seen this in practice by studying gravitational and electrical potential energy in past courses.

Gravitational Potential Energy

– An object held above a table contains gravitational potential energy. When released, it will fall towards the table in order to minimize its potential energy and release it as kinetic energy.

Electrical Potential Energy

– Negative charge accumulating at the anode of a battery contains electrical potential energy. Charge will move through a circuit towards the cathode in order to minimize its potential energy and release it as heat, light, etc.

Potential energy, in general, is a concept involving the potential to perform an action before that action has actually been performed. Therefore, it is inversely related to stability. Just as an object is more stable on a surface than when suspended in the air, an object with high potential energy experiences a great deal of instability before releasing that energy into something productive and becoming more stable as a result.

These same principles apply to bonding in chemistry.

Energy and Stability – Ionic Bonding

The potential energy released when ionic bonds form is known as lattice energy. The higher the lattice energy, the stronger the bond will be between a pair of ions in an ionic compound. Similarly, we can refer to the potential energy released when an ionic compound dissociates in solution with water as hydration energy.

Based on the principle of minimum energy, whichever form of energy results in a larger release of potential energy will be the state in which the compound is most stable. If hydration energy is higher, the compound will likely dissociate and form an aqueous solution. If lattice energy is higher, the compound is less likely to be soluble.

A chemical reaction will occur if the total potential energy of the products is lower than the potential energy of the reactants. The net ionic equation summarizes the compound for which the potential energy is changing.

Example 1:

Complete Chemical Equation:AgNO3 (aq) + KCl (aq) → AgCl (s) + KNO3 (aq)Net Ionic Equation:Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

Lattice Energy of silver chloride > Hydration Energy of silver chlorideTherefore, AgCl will form a precipitate instead of remaining in solution.

Example 2:

Complete Chemical Equation:FeSO4 (aq) + 2 KI (aq) → FeI2 (aq) + K2SO4 (aq)Net Ionic Equation:No net ionic equation (all compounds are aqueous)

Hydration Energy of all compounds > Lattice Energy of all compoundsTherefore, all compounds will remain in solution and no reaction will occur.

Energy and Stability – Covalent Bonding

Similar to ionic bonding, molecular compounds will also tend towards their lowest state of potential energy. As the atoms move closer to one another and their nuclei begin to attract the shared electrons, their potential energy decreases.

The added wrinkle to this type of bonding lies in bond length (represented on the graph as Internuclear distance). At a certain point, the atoms reach an optimal separation at which they can share their electrons. Moving beyond this point (shortening the length of the bond) increases the potential energy needing to be released due to the fact that the positively-charged nuclei are strongly repelling one another.

Review Questions

For the following ionic compounds, state which type of energy you expect to be higher (lattice or hydration). Hint: Refer to the examples in the note to help you.

• Silver chloride: Lattice Energy

• Potassium iodide: Hydration Energy

• Iron(II) sulfate: Hydration Energy

Outline the relationship between potential energy and stability.

• Potential energy and stability are inversely related, meaning that higher potential energy results in lower stability.

Extension

Interestingly, internuclear distance is also relevant to the strength of an ionic bond. What characteristic of an atom (often referenced as a trend in the periodic table) might impact the internuclear distance between two ions?

• Atomic/Ionic Radius

Extension

Based on the principles described in this note, do you think increasing the number of bonds (i.e. double bonds or triple bonds) would increase or decrease the bond length? Why?

• As two atoms get closer to one another, they attract each other’s electrons in order to complete their valence shell. This continues to an optimum point which will be a certain length. As a result, it is logical to conclude that if more electrons are shared, the atoms will be closer to one another when they achieve minimum potential energy. This results in double and triple bonds being shorter than single bonds in a molecule.

Chemical Bonding and Electronegativity

In the atomic tug-of-war game, are all the atoms equally strong? Now that you are familiar with ionic and covalent bonds, we will look at them a little more closely. Evidence shows that the atoms of some elements attract electrons more strongly than do the atoms of other elements.

Electronegativity

In Chapter 1, you learned that several atomic properties follow periodic trends. As you now explore chemical bonding in more detail, it is time to learn about another atomic property: electronegativity. Electronegativity is the ability of an individual atom, when bonded, to attract bonding electrons to itself. The American chemist Linus Pauling (1901–1994) developed the concept of electronegativity in 1922. He also proposed a scale of electronegativity values. Simply put, comparing the electronegativity values of two atoms indicates the likelihood of those two atoms being part of an ionic or a molecular compound.

Pauling assigned the highest electronegativity value to fluorine (4.0) and the lowest value to francium (0.7). Fluorine is the element that has the greatest ability to attract a bonding pair of electrons. Francium has the lowest ability to attract bonding electron pairs to itself. In a way, electronegativity indicates how strong an atom is in a tug-of-war over bonding electrons. An element with a high electronegativity is very good at pulling a pair of electrons toward itself, whereas an element with a low electronegativity is not.

Electronegativity Trends

In general, the electronegativity of the elements increases from left to right on the periodic table and decreases from top to bottom. To explain this trend, it helps to consider atomic radius. Remember that atomic radius decreases as you move from left to right across the periodic table. As the size of an atom increases, the attraction it has for a shared electron pair weakens. Therefore, electronegativity follows the opposite trend to atomic radius on the periodic table. As atomic radius increases, electronegativity decreases and vice versa.

The electronegativity of an element cannot be measured experimentally. Instead, it is calculated for each element using physical properties such as ionization energy. Accepted values of electronegativity are given in the periodic table at the back of this textbook. Values for some of the more common elements are also given below.

Determining Whether a Compound Is Ionic or Molecular

General rule for ionic compounds:

• Form when a metal combines with a non-metal

General rule for molecular compounds:

• Form when two or more non-metals combine

To predict whether a bond is ionic or covalent, we must calculate the difference in the elements’ electronegativities. This quantity is called the electronegativity difference (symbol ΔEN).

The greater the electronegativity difference, the more likely it is that the bond is ionic.

Consider the electronegativity difference between sodium and chlorine. If we subtract the smaller value (0.9) from the larger value (3.2), we get:Electronegativity difference for Na-Cl = 3.2 – 0.9 = 2.3.

The relatively large difference indicates that chlorine is strong enough to take an electron from sodium, resulting in a chloride anion (Cl⁻) and a sodium cation (Na⁺). The resulting bond would be ionic.

Covalent Bonds: Non-Polar and Polar

When atoms are identical (ΔEN = 0), they share electrons equally. This type of covalent bond is called a non-polar covalent bond. Non-polar covalent bonds form for all diatomic elements.

Example: Cl2(g)

Polar Covalent Bonds

When two covalently bonded atoms have an electronegativity difference between 0 and 1.7, they do not share electrons equally. Consider the electronegativity difference between hydrogen (2.2) and chlorine (3.2). Chlorine has a higher electronegativity value and will attract electrons away from the lower-valued hydrogen. Therefore, the chlorine atom will have a more negative charge overall than the hydrogen atom, leading to a formation of a polar covalent bond.

Review

1. Explain how the trends of electronegativity in the periodic table relate to those of atomic radius.

2. Using only their relative position on the periodic table, arrange the following elements in order of increasing electronegativity: K, Cs, Br, Fe, Ca, F, Cl

3. Describe how we can use electronegativity values to predict the types of bonds that will form in a compound.

4. Distinguish between a polar covalent bond and a non-polar covalent bond.

5. For each bond listed, determine the electronegativity difference and predict what type of bond (non-polar covalent, polar covalent, or ionic) that would form between the two elements:(a) Ca-S(b) H-F(c) P-H(d) C-Cl(e) C-O(f) Li-Cl

6. Identify the more polar bond in each of the following pairs:(a) H-F and H-Cl(b) O-H and C-H(c) C-N and N-N

7. Predict whether the chemical bonds in each of the following compounds will be non-polar covalent, polar covalent, or ionic:(a) N2(g)(b) NH3(g)(c) H2O(l)(d) FeO(s)(e) MgCl2(s)

Polar Bonds and Polar Molecules

Recall that a covalent bond is either polar or non-polar. A polar covalent bond is formed when atoms with very different electronegativities result in an unequal sharing of the bonding electron pair. In a polar covalent bond, the shared electrons tend to be closer to the atom with the higher electronegativity. The electron density is higher around the more electronegative atom, resulting in a negatively charged pole, and less electron density around the less electronegative atom, resulting in a positively charged pole.

There are varying degrees of difference in electronegativity. The extent to which a bond is polar depends on that electronegativity difference. Consider the following examples:

• H-H: a non-polar covalent bond

• N-H: a polar covalent bond

• O-H: a very polar covalent bond

Molecular Polarity

In your own words, define:Polar molecule: a molecule with an uneven distribution of electrons that results in a positive charge at one end and a negative charge at the other end.

Non-polar molecule: a molecule with equally distributed electrons between atoms that results in no localized charges.

Polar liquids, such as water, have some special physical properties. You may have noticed that water will “stick” to itself in the form of droplets, or that you can cause a stream of water to move towards a negatively charged object. What might be the cause of these properties?

Think about how the charges of each molecule of water must line up.

Polar Molecules

We know that water molecules are polar. Why?

• Water is made up of one oxygen and two hydrogen atoms

• There are two O-H bonds

• The ΔEN between O and H is 1.4, so each bond is polar covalent

• Oxygen attracts the bonding electrons more strongly, so there are negative charges around the oxygen atom and positive charges around the hydrogen atoms

• Water is bent and has TWO ends: a negatively charged oxygen end and a positively charged hydrogen end.

Non-Polar Molecules

Let’s now consider oxygen as a diatomic molecule:

• O2 contains two of the same atom (same electronegativities)

• Both covalent bonds are non-polar; therefore, the whole molecule is non-polar

• Diatomic molecules made up of identical atoms are ALWAYS non-polar

What about a polyatomic molecule? Can it be non-polar? Let’s look at carbon dioxide.

• CO2 is made of 3 atoms joined by 4 covalent bonds: 2 double-bonds each between oxygen and carbon

• Oxygen has the higher electronegativity (ΔEN = 1), so each bond is polar

• BUT! The two polar covalent bonds are arranged symmetrically across the central carbon atom; therefore, the positive and negative charges will cancel out

• The molecule is NON-POLAR overall.

How to determine if a molecule is polar or non-polar

To determine if a molecule is polar, we must consider BOTH the types of bonds (polar or non-polar) AND the arrangement of those bonds around the central atom in the molecule.

Intermolecular Forces Summary

Summary of Non-Covalent Interactions:

Problems:

1. Name the dominant (strongest) intermolecular force in the following pairs:

   • Methane and Methane (CH4): London Forces

   • Ethanol and Ethanol: Hydrogen Bonding

   • Water and Water: Hydrogen Bonding

   • NH3 and NH3: Hydrogen Bonding

   • Cyclohexanone and Cyclohexanone: Dipole-Dipole

   • Cyclohexanol and Cyclohexanol: Hydrogen Bonding

   • HCl and HCl: Dipole-Dipole

   • CO2 and CO2: London Forces

   • CCl4 and CCl4: London Forces

   • CH2Cl2 and CH2Cl2: Dipole-Dipole

2. If the pairs of substances listed below were mixed together, list the non-covalent interaction(s) that are involved:

   • NH3 and H2O: A, B, C (Hydrogen Bonding, Dipole-Dipole, London Forces)

   • Mg2+ and H2O: D (Ion-Dipole)

   • Cl2 and H2: C (London Forces)

   • Acetic acid and H2O: A, B, C (Hydrogen Bonding, Dipole-Dipole, London Forces)

   • SO2 and H2O: A, B, C (Hydrogen Bonding, Dipole-Dipole, London Forces)

   • SO2 and H2S: B, C (Dipole-Dipole, London Forces)

   • Ethane (CH3CH3) and Methane (CH4): C (London Forces)