Electronegativity and Valency (HSC SSCE Chemistry): Revision Notes
Electronegativity and Valency
Understanding electronegativity
Electronegativity is a measure of how strongly an atom attracts electrons towards itself in a chemical bond. You can think of it as the "electron-pulling power" of an element. This property was introduced in earlier sections and plays a crucial role in determining what type of chemical bond will form between two elements.
When two atoms bond, they may either share electrons (forming a covalent bond) or transfer electrons from one atom to another (forming an ionic bond). The difference in electronegativity between the two atoms helps us predict which type of bonding will occur.
The greater the difference in electronegativity between two atoms, the more likely they are to form an ionic bond through electron transfer rather than a covalent bond through electron sharing.
Predicting ionic versus covalent bonding
The 1.5 rule
There is a useful guideline for predicting whether two elements will form an ionic or covalent compound based on their electronegativity difference:
- If the electronegativity difference is greater than 1.5, the compound will be ionic
- If the electronegativity difference is less than 1.5, the compound will be covalent
Why this rule works
For an ionic compound to form, one atom must have a much stronger attraction for electrons than the other. This allows for the complete transfer of electrons from one atom to another. If both atoms have similar attractions for electrons, they are more likely to share electrons equally, forming a covalent bond instead.
Worked examples
Let's look at some examples to see how this rule works in practice:
Worked Example 1: Hydrogen and sulfur
Could hydrogen and sulfur form H₂S as an ionic or covalent compound?
- Electronegativity of H = 2.20
- Electronegativity of S = 2.58
- Difference = 2.58 - 2.20 = 0.38
Since 0.38 is much less than 1.5, hydrogen sulfide (H₂S) is a covalent compound. The two hydrogen atoms each share a pair of electrons with the sulfur atom.
Worked Example 2: Calcium and nitrogen
- Electronegativity of Ca = 1.00
- Electronegativity of N = 3.04
- Difference = 3.04 - 1.00 = 2.04
Since 2.04 is greater than 1.5, the compound between calcium and nitrogen is ionic.
Worked Example 3: Phosphorus and chlorine
- Electronegativity of P = 2.19
- Electronegativity of Cl = 3.16
- Difference = 3.16 - 2.19 = 0.97
Since 0.97 is less than 1.5, the compound between phosphorus and chlorine is covalent.
Best practice for prediction
While the 1.5 rule is helpful, the most reliable method for determining bond type is to consider the electron requirements of each element:
- If one element needs to lose electrons while the other needs to gain electrons, the compound will be ionic
- If both elements need to gain electrons, the compound will be covalent
This approach considers the fundamental nature of how atoms achieve stable electron configurations.
Valency and the periodic table
What is valency?
Valency (also called valence) is the combining capacity of an element—it tells us how many chemical bonds an atom can form. Understanding valency is essential for predicting chemical formulas and understanding how elements combine.
Valency rules for main-group elements
The periodic table provides a systematic way to determine the common valency of main-group elements. Here are the key rules:
For Groups 1 and 2:
- Valency = the group number
- Group 1 elements (Li, Na, K, Rb, Cs) have valency of 1
- Group 2 elements (Be, Mg, Ca, Sr, Ba) have valency of 2
For Group 13:
- Valency = group number minus 10
- This gives a valency of 3 for elements like B and Al
For Groups 14 to 17:
- Valency = 18 minus the group number
- Group 14: valency = 18 - 14 = 4
- Group 15: valency = 18 - 15 = 3
- Group 16: valency = 18 - 16 = 2
- Group 17: valency = 18 - 17 = 1
For Group 18:
- Valency = 0 (noble gases already have stable electron configurations)
Why these rules work
These valency patterns arise from how atoms achieve stable electron configurations:
Understanding the Electron Requirements:
- For Groups 1, 2, and 13: The valency represents the number of electrons the atom needs to lose to achieve a noble gas configuration
- For Groups 14 to 17: The valency represents the number of electrons the atom needs to gain to achieve a noble gas configuration
- For Group 18: These elements already have stable configurations, so they don't need to form bonds (valency = 0)
Common valencies by group
The following table summarises the common valencies of elements in the main groups of the periodic table and the types of compounds they typically form:
| Group | Elements | Common valency | Type of compound |
|---|---|---|---|
| 1 | Li, Na, K, Rb, Cs | 1 | Ionic |
| 2 | Be, Mg, Ca, Sr, Ba | 2 | Ionic |
| 13 | B, Al | 3 | B forms covalent; Al forms ionic |
| 14 | C, Si | 4 | Covalent |
| 14 | Ge, Sn, Pb | 2 and 4 | Ionic with valency 2; can be covalent with valency 4 |
| 15 | N, P | 3 or 5 | Covalent |
| 16 | O | 2 | Ionic or covalent |
| 16 | S | 2, 4, 6 | Ionic (valency 2 only) or covalent |
| 17 | F, Cl, Br, I | 1 | Ionic or covalent |
Important notes about the table
Key observations from the valency table:
- Group 1 and 2 elements typically form ionic compounds because they readily lose electrons
- Non-metals (Groups 14-17) can form either ionic or covalent compounds depending on what they bond with
- Carbon and silicon (Group 14) primarily form covalent compounds
- Oxygen and sulfur (Group 16) can form both types of compounds, with oxygen commonly showing valency of 2
- Halogens (Group 17) can form either ionic compounds (with metals) or covalent compounds (with non-metals)
Transition metals and valency
Unlike main-group elements, you cannot deduce the valency of transition metals from their position in the periodic table. However, there are some useful patterns:
Transition Metal Patterns:
Most transition metals form M²⁺ cations (where M represents the metal). Many also form cations with other charges.
Examples include:
- Copper: Cu⁺ and Cu²⁺
- Iron: Fe²⁺ and Fe³⁺
- Chromium: Cr²⁺ and Cr³⁺
Exceptions:
- Zinc forms only Zn²⁺
- Silver mainly forms Ag⁺
Simple compounds of transition metals are typically ionic in nature.
Remember!
Key Points to Remember:
- Electronegativity measures how strongly an atom attracts electrons in a bond—it's the atom's "electron-pulling power"
- Use the 1.5 rule to predict bonding: if the electronegativity difference is greater than 1.5, the compound is ionic; if less than 1.5, it's covalent
- Valency indicates the combining capacity of an element—how many bonds it can form
- For main-group elements, you can determine valency from the periodic table: Groups 1-2 use the group number; Groups 14-17 use 18 minus the group number
- The valency rules work because they reflect the number of electrons atoms need to gain or lose to achieve a stable noble gas configuration