Ionic and Covalent Compounds

Drawing Ionic and Covalent Compounds

• The correct Lewis structure for the ionic compound formed between sodium (Na) and chlorine (Cl) is Na⁺ [Cl]⁻.
• A covalent bond forms between two non-metal atoms in a covalent compound.
• N₂ has a triple bond.
• When drawing the Lewis structure of a molecule, one rule to follow is that all atoms must have a complete outer electron shell.
• The overall charge of an ionic compound is neutral.
• Magnesium chloride (MgCl₂) is an ionic compound formed through the transfer of electrons from magnesium to chlorine, resulting in Mg²⁺ and two Cl⁻ ions.
• Ionic compounds conduct electricity when dissolved in water due to the mobility of ions, whereas covalent compounds generally do not conduct electricity under any conditions unless they undergo ionization or have free-moving electrons.

Properties of Ionic, Covalent, and Metallic Compounds

• High solubility in water is a property typical of ionic compounds.
• Metallic bonds are best described as electrons pooled and free to move around.
• Covalent compounds are poor conductors of electricity in all states.
• Copper (Cu) would be classified as a metallic compound.
• Ionic compounds tend to have high melting points because the bonds between ions are very strong and require significant energy to break.
• Ionic compounds like sodium chloride conduct electricity when dissolved in water because the ions are free to move and carry charge, but they do not conduct electricity in their solid state because the ions are held in fixed positions within the crystal lattice.
• Metallic copper exhibits malleability and conductivity due to its structure of positive ions immersed in a "sea" of electrons, allowing electrons to move freely and ions to slide past each other without breaking bonds, whereas sulfur, a covalent network solid, is brittle due to its directional bonds that break under stress.

History of the Atomic Model

• J.J. Thomson is credited with the discovery of the electron.
• Bohr’s model of the atom proposed that electrons orbit the nucleus in fixed paths.
• Ernest Rutherford is most associated with the “gold foil” experiment.
• The modern atomic model is based on the principles of electron clouds.
• A correct statement about Dalton’s atomic theory that atoms are indivisible and indestructible.
• Rutherford's model introduced the concept of a small, dense, positively charged nucleus surrounded by empty space, while Thomson's model proposed a "plum pudding" structure with electrons dispersed throughout a positive matrix. These discoveries challenged the earlier view of the atom as an indivisible, uniform sphere.
• The concept of the atom has evolved from Dalton's indivisible sphere to the quantum mechanical model through a series of experimental findings, including Thomson's discovery of electrons, Rutherford's nuclear model, and Bohr's quantized orbits, each leading to refinements based on new evidence and a progressively more accurate understanding of atomic structure.

Periodic Trends

• As you move across a period from left to right on the periodic table, the electronegativity of elements increases.
• Fluorine (F) has the highest electronegativity.
• Atomic radius generally increases as you move from top to bottom down a group.
• As you move from top to bottom in a group on the periodic table atomic radius increases.
• The ionization energy of an element is defined as the energy needed to remove an electron from a neutral atom.
• Fluorine has a higher electronegativity than iodine because fluorine is smaller, resulting in a greater effective nuclear charge that attracts bonding electrons more strongly.
• Ionization energy increases across a period due to increasing nuclear charge and decreases down a group due to increased electron shielding; these trends relate to the number of protons and the electron shielding effect by affecting how strongly the nucleus attracts the outermost electrons.

Isotope Calculations, Definition, and Uses

• An isotope of an element has the same number of protons but different numbers of neutrons.
• The isotope carbon-14 is used in radiocarbon dating.
• The atomic mass of an element is the average mass of all its isotopes, weighted by abundance.
• Carbon-12 and Carbon-14 are examples of an isotope.
• Isotopes differ in their number of neutrons.
• With Isotope A having a mass of 10.0 amu and an abundance of 60%, and Isotope B having a mass of 11.0 amu and an abundance of 40%, the average atomic mass calculates to 10.4 amu by multiplying each isotope's mass by its abundance and summing the results.
• Isotopes are used in medicine for diagnostic imaging and therapeutic treatments, such as radioactive iodine for thyroid disorders. While these applications offer benefits like precise targeting, they also carry risks associated with radiation exposure, including potential harm to healthy tissue.

Polar and Non-Polar Covalent Compounds

• A molecule that is polar will have a net dipole moment.
• Cl-Cl bond is most likely to be non-polar.
• A polar covalent bond occurs when the electronegativity difference between two atoms is significant.
• CO₂ is non-polar despite having polar covalent bonds
• In H₂O, the dipole-dipole forces will be the strongest.
• CO₂, despite having polar bonds, is a non-polar molecule due to its linear geometry, which cancels out the bond dipoles, resulting in no net dipole moment.
• A non-polar substance will not dissolve well in water due to the lack of attraction between non-polar molecules and polar water molecules; non-polar substances prefer to interact with other non-polar substances.