Ions, Atomic Structure and Prefixes

• Prefixes for naming chemical compounds indicate the number of atoms: mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca

Atomic Structure and Ions

• The periodic table can determine an atom's structure and the charge when it forms ions.
• Subatomic particles in an atom include protons (positive charge, atomic number, mass of 1), neutrons (no charge, mass of 1, found in the nucleus), and electrons (negative charge, orbits around the nucleus).
• Neutral atoms have no charge, with the same number of protons and electrons.
• The period number indicates the number of electron shells.
• The group number indicates valency or the number of valence electrons, corresponding to the charge of elements as ions.
  • Valency is the number of bonds an atom can form, or the number of atoms gained/lost/shared to fill its outer shell.
  • Valence electrons participate in bonding.
• Net charge is the overall electric charge an atom possesses.
• Molecules consist of two or more atoms joined by a covalent bond.
  • Each atom has a full outer electron shell and is stable.
• Diatomic elements include hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, and iodine; and they exist in pairs.

Formation of Ions and Ionic Bonds

• Cations are positive ions that lose electrons.
  • Ionization energy (I.E.) is the energy needed to remove an electron from an atom.
  • Example: Na(g) -> Na+(g) + e-
• Anions are negative ions that gain electrons and are larger due to electron repulsion in the outer shell.
  • Example: Cl(g) + e- -> Cl-
• Full outer shells have the electronic configuration of a noble gas.
• Ionic bonds are a metal and non-metal compound in a ratio allowing each atom to achieve a noble gas electron configuration.
  • Electrons are transferred.
  • Metals and nonmetals form neutral salts with an overall charge of zero, belonging to groups 1, 2, 6, 7 (sometimes 3).
• Compounds:
  • K+ has a valency of +1 by losing an electron.
  • S-2 has a valency of -2 by gaining two electrons.
  • K2S is formed by two potassium ions and one sulfide ion, resulting in a neutral compound/salt.
• The charge for an ionic compound is overall zero.
• Naming ionic compounds involves stating the cation first, then the anion, without writing the + or - charges.
• Roman numerals are used to indicate the charge in the worded chemical name, such as copper (II) oxide.
• The general formula for ionic compounds is derived by swapping and dropping charges to determine the subscript for each ion.
• The superscript number indicates the charge on the ion, while the subscript number indicates the number of ions of that type in the compound
• Polyatomic ions include:
  • CO3^2- (carbonate)
  • SO4^2- (sulfate)
  • PO4^3- (phosphate)
  • OH- (hydroxide)
  • NO3- (nitrate)
  • NH4+ (ammonium)

Covalent Bonds

• Covalent bonds form when two non-metal atoms share a pair of electrons.
  • There are no charges involved.
  • Prefixes such as mono, di, and tri are used.
• An ionic bond results from the attraction of positive and negative ions, where the positive component is a metal, and the negative is a non-metal.

Classifying Compounds (Hydrocarbons)

• Organic chemistry involves compounds containing carbon which are present in all living and many non-living things.
  • Carbon has four valence electrons which allow it to form four bonds.
  • Carbon bonds very well to other carbon atoms which leads to long chains and cyclic molecules and readily bonds to other elements like hydrogen, oxygen, nitrogen, sulfur, and chlorine.
• Families are classifications of carbon compounds with common characteristics, known as homologous series that share the same ratio.
  • Hydrocarbons only contain hydrogen and carbons.
  • Homologous series have the same general formula, with each member different by one CH2 group with similar chemical properties and a gradual change in physical properties as more carbon atoms are added.
  • Examples include petrol, DNA, and sugar.

Homologous Series of Hydrocarbons

• The classifications within homologous series differ in physical properties.
• Alkanes, Alkenes, and Alkynes
  • Alkanes (CnH2n+2) are the simplest hydrocarbons with single bonds between carbon atoms.
  • Alkenes (CnH2n) contain double bonds between carbon atoms and are unsaturated hydrocarbons.
  • Alkynes (CnH2n-2) contain triple bonds between carbon atoms and are unsaturated hydrocarbons.
• Saturated compounds have the maximum number of hydrogen atoms attached to the carbon atoms and are typically found in alkanes.
• Unsaturated compounds have fewer hydrogen atoms because of the presence of double or triple bonds found in alkenes and alkynes.

Examples of Hydrocarbons

• Alkanes: methane (CH4), ethane (C2H6), propane (C3H8)
• Alkenes: ethene (C2H4), propene (C3H6)
• Alkynes: ethyne (C2H2) or acetylene.

Crude Oil

• Crude oil is made of the remains of dead sea animals and requires heat and pressure to form.
• Technology used to find it includes seismic surveys, gravitational surveys, geological mapping, and technology for use in plastic and vehicles.

Types of Chemical bonds

• Ionic: formed when a metal reacts with a nonmetal, held together by electrostatic forces (ex. NaCl)
• Covalent: formed when two nonmetals react, held together by the electrons that are shared (CO2)
• held together by metallic bonds (Fe)

Other chemical types and examples are as follows:

• Acidic: release hydrogen ions when dissolved in water, pH is less than 7 (HCl)
• Basic: release hydroxide ions when dissolved in water, pH is greater than 7 (NaOH)
• Organic: include hydrogen, oxygen, nitrogen (CH4)
• Inorganic: do not contain carbon-hydrogen bonds (NaCl)

IUPAC

• The International Union of Pure and Applied Chemistry (IUPAC) standardizes chemical naming, terminology, and measurements, which prevents errors/miscommunication and improves real world applications.

Chemical Reactions:

• Expressed using equations which should then be balanced.
• Reactants are transformed to products (e.g., Na + Cl -> NaCl).
• Signs of a chemical reaction include color changes, temperature changes, giving off light/sound, producing gas or forming a solid.
• State of chemicals can be solid (s), liquid (l), gas (g), or aqueous (aq) when dissolved in a solution.

Exothermic Reactions:

• Release energy and can occur spontaneously.
• Form of combustion, respiration, neutralization, or any reaction of metal with acid.
• Reactants have more energy than the products, resulting in the release of energy.

Combustion:

• An exothermic process where heat and light are produced, which requires heat, oxygen, and fuel with the following format: Hydrocarbon + oxygen -> Water + Carbon Dioxide

Energy Transfer:

• In reactions between magnesium and hydrochloric acid, no external heat is needed, so during the reaction chemical energy in reactants convert to thermal energy, which then disperses and returns to normal.
• Endothermic reactions absorb energy and decrease in temperature, with thermal decomposition, photosynthesis, electrolysis all requiring positive. In these reactions the products have more energy than the reactants.
• In reactions between ammonium nitrate and water, thermal energy converts to chemical energy, causing the temperature to fall.

Endothermic and Exothermic Differences

• Exothermic reactions release energy from system .
• Endothermic reactions absorbs energy, as heat.
• Most reactions are exothermic except gas to solid.
• In order from solid-liquid-gas it´s endothermic because they absorb heat.
• In order from gas-liquid-solid it´s exothermic because they release heat.