Atomic Structure & Bonding: Electron Config, Isotopes & Bonds

Two shells indicate Period 2 element. Six outer electrons correspond to Group 16. Oxygen fits both conditions: atomic number 8 gives configuration 2,6. Fluorine has seven valence electrons, carbon has four, chlorine has three shells. Therefore oxygen uniquely satisfies shell count and valence electron requirement, confirming identification logically and structurally.

Beryllium forms +2 ion while fluorine forms −1 ion. To achieve charge neutrality, total positive and negative charges must balance. One Be²⁺ requires two F⁻ ions since 2 × (−1) = −2 balances +2. Ratio becomes 1:2. Therefore empirical formula is BeF₂. Any other combination would create net charge, violating electrical neutrality principle in ionic compounds.

Be²⁺ carries +2 charge and N³⁻ carries −3 charge. To balance, use lowest common multiple of 6. Three Be²⁺ ions give +6 total. Two N³⁻ ions give −6 total. Ratio becomes 3:2. Therefore empirical formula is Be₃N₂. Any smaller ratio fails to neutralize total charge completely in ionic lattice.

Metallic bond strength depends on number of delocalized valence electrons. More electrons increase electrostatic attraction between positive metal ions and electron sea. Stronger attraction raises melting point and hardness. Fewer electrons weaken bonding. Neutrons do not influence bonding strength significantly. Therefore increased delocalized electron density directly enhances metallic bonding strength quantitatively.

Silicon has atomic number 14, meaning 14 electrons in neutral form. Shell filling follows energy levels: first shell holds 2 electrons, second holds 8. That accounts for 10 electrons. Remaining 4 electrons occupy third shell. Therefore distribution becomes 2,8,4. This arrangement explains silicon’s four valence electrons, giving it semiconductor properties and tetravalent bonding behavior in compounds.

Carbon requires four shared electrons, oxygen requires two, and each hydrogen requires one. Carbon forms two single bonds with hydrogen atoms, using two electrons. Remaining two electrons form double bond with oxygen, satisfying octet rule for carbon and oxygen. Hydrogen achieves duet rule. This structure matches valency requirements and observed molecular geometry experimentally.

Group 13 Period 3 element is aluminum. Aluminum has atomic number 13. Electron distribution becomes 2 in first shell, 8 in second, and 3 in third, totaling 13 electrons. Therefore configuration 2,8,3 is correct. This explains aluminum forming +3 ions by losing three valence electrons during bonding reactions commonly observed.

Ionic compounds form giant electrostatic lattices with strong Coulombic attractions between oppositely charged ions. These forces extend in all directions, requiring large energy to overcome. Metallic bonding strength varies widely, and simple molecular covalent substances have weaker intermolecular forces. Therefore ionic solids typically exhibit very high melting points compared with most covalent or metallic substances.

Weighted average uses fractional abundances. Convert percentages: 0.005%, 0.720%, 99.275%. Multiply masses: 234×0.00005 = 0.0117; 235×0.00720 = 1.692; 238×0.99275 = 236.2695. Adding gives 237.9732. Rounding to four decimal places remains 237.9732. Large abundance of U-238 dominates calculation significantly, explaining value proximity to 238 atomic mass units.

Ionic compounds consist of oppositely charged ions arranged in a rigid lattice. Strong electrostatic attractions hold ions in fixed positions, making the structure brittle when layers shift and repel. Electricity conduction requires free-moving charged particles. In solid state, ions are fixed and cannot move. When molten or dissolved, ions become mobile, allowing charge flow. This explains brittleness and conductivity only in liquid or aqueous states.

Malleability arises from metallic bonding structure. Metal atoms form positive ions in lattice surrounded by delocalized electrons. When force is applied, layers slide without breaking bonds because electron sea redistributes charge evenly. This prevents fracture. In contrast, ionic solids shatter due to repulsion when layers shift. Therefore metallic bonding structure directly explains malleability and shaping ability.

Alkali metals occupy Group 1. Each atom has one valence electron, which is easily lost to form +1 ions. Low ionization energy explains extreme reactivity. Their metallic bonding is relatively weak due to single valence electron, making them soft and silvery. Reactivity increases down group as atomic radius increases and outer electron attraction decreases significantly.

Isotopes differ in neutron number while proton number remains constant. Changing neutrons alters atomic mass but not chemical identity or valence electrons. Since melting point slightly depends on mass differences affecting atomic vibrations, small variations may occur. However proton number and electron configuration remain unchanged. Therefore relative atomic mass differs between isotopes fundamentally.

Relative atomic mass equals weighted average. Multiply each isotope mass by its fractional abundance: 19×0.20 = 3.8, 20×0.09 = 1.8, 21×0.52 = 10.92, 22×0.19 = 4.18. Adding gives 3.8 + 1.8 + 10.92 + 4.18 = 20.7. Since percentages total 100%, no further division needed. Therefore, weighted average mass equals 20.7 atomic mass units accurately.

Strontium belongs to Group 2 and is positioned further left than vanadium. Reactivity of metals increases moving left across a period and down a group due to easier electron loss. Strontium loses two electrons readily, making it highly reactive. Vanadium is transition metal with partially filled d orbitals and lower reactivity. Thus statement claiming vanadium more reactive is incorrect.