Chapter 2 Notes
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This 9 page Class Notes was uploaded by Shannan Dillen on Thursday September 29, 2016. The Class Notes belongs to CHEM 1510 at Ohio University taught by Shadrick Paris in Fall 2016. Since its upload, it has received 3 views. For similar materials see Fundamentals of Chemistry I in Chemistry at Ohio University.
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Date Created: 09/29/16
Chapter 2: The Components of Matter Element – simplest type of substance with unique physical and chemical properties -consists of only one type of atom -cannot be broken down into simpler substances Molecule – structure that consists of 2 or more atoms that are bonded together -chemically and covalently bonded together -behaves as an independent unit Compound – 2 or more elements that are chemically combined -fixed, specific ratio by mass Mixture – group of 2 or more substances (could be elements and compounds) -physically intermingled -components can be varied by mass Mixture – things can be separated by using the physical properties of the elements Compound – chemically combined; cannot by physically separated Mixtures Heterogeneous – one or more visible boundaries Ex: oil and water Homogeneous – can also be called a solution; mixed at atomic level Solutions in water – aqueous solutions A mixture can be separated into its component substances (elements and/or compounds) by physical methods. Law of Mass Conservation – the total mass of a substances does not change during a chemical reaction Law of Definite (Constant) Composition – no matter the source, a compound is composed of the same elements in the same parts (fractions) by mass Law of Multiple Proportions – if elements A and B react to form 2 different compounds, the different masses of B that combine with a fixed mass of A can be expressed as a ratio of small whole numbers Dalton’s Atomic Theory 1. All matter consists of atoms. 2. Atoms of one element cannot be converted into atoms of another element. 3. Atoms of an element are identical in mass and other properties and are different from the atoms of any other element. 4. Compounds result from the chemical combination of a specific ratio of atoms of different elements. Robert Millikan, 1909 -charge of electrons is -1.602 * 10 -19C (coulomb) -mass of electrons is 9.109 * 10 -28g Ernest Rutherford, 1909 -the atom is mostly space -there’s a dense region in center, dubbed the nucleus -the nucleus contains all of the positive charge Atom is electrically neutral. (charge=0) The atom nucleus consists of protons and neutrons. “Diameter” of an atom is ~20,000 times greater than that of its nucleus. Protons (p) +1 Neutron (n) 0 - Electrons (e ) -1 X = atomic symbol of the element Z = atomic number (different for every element) A = mass number Charge = net charge due to number of positive charges and negative charges WHOLE NUMBERS # of protons (z) = defines which ELEMENT # of electrons (z – charge) = defines what CHARGE Charge = 0 – ATOM Charge does not = 0 – ION # of neutrons (N = A – z) = defines MASS Isotopes – atoms of an element with same number of protons, but different number of neutrons Isotopes have the same atomic number, but a different mass number. (Note: N = A – z) For any specific element (and atomic number Z), there can be a multitude of isotopes. Some synthetic (man-made); some are naturally occurring Some isotopes are stable; others unstable and undergo nuclear decay Atomic and Isotopic Mass Atoms are TINY – doesn’t make sense to use standard mass units Individual atomic and isotopic masses are measured relative to an atomic standard called the atomic mass unit (amu). 1 amu = 1/12 the mass of one atom of carbon-12 (EXACT) Why carbon-12? -has 6 protons, 6 neutrons, 6 electrons 1 amu = mass of one proton and one neutron 1 amu = 1.66054 * 10 -2g Each isotope will have its own isotopic mass, in units of amu, which will be close to its mass number, but not an exact whole number. The ONLY isotope whose isotopic mass is an exact, whole number is carbon-12. For elements, weighted average atomic mass can be calculated from isotope masses of all its naturally occurring, stable isotopes, taking into account each isotope’s respective natural abundance. Atomic mass = (isotopic mass A)(%A/100) + (isotopic mass B)(%B/100) + … The atomic mass will tend to be closest in value to that of the most abundance. The Periodic Table Organized by increasing atomic mass from left to right and from top to bottom Rows – periods Columns – groups Groups are split into A and B and are each numbered 1 through 8 Table grouped and divided based on similar physical and chemical properties of the elements Main group elements – A groups Transition metals/elements – B groups Rare earth metals – inner transition elements Some groups (or “families”) have been given names - Alkali Metals (Group 1A) H not counted - Alkaline Earth Metals (Group 2A) - Halogens (Group 7A) - Noble Gases (Group 8A) Elements in same group tend to have similar chemical properties. Metals - Solids at room temperature o Hg is liquid at room temperature - Lustrous (“shiny”) - Conduct both heat and electricity well 3 - High densities (anywhere from around 1 to 20 g/cm ) - Malleable and ductile Nonmetals - Solids at room temperature o C, P, S, Se, I - Liquid at room temperature o Br - Gases at room temperature + o H , N, O, F, Cl, noble gases (group 8A) - “dull” - Conduct both heat and electricity poorly - Low densities - Brittle Metalloids B, Si, As, Ge, Sb, Te - Solids at room temperature - Properties “between” those of metals and nonmetals Polyatomic Elements Some elements occur in nature as molecules Ex: P4, S8, C60 Specific group of two – diatomic elements H 2 N 2 O 2 F2, Cl2, B2 ,2I -only diatomic in pure, elemental forms “Br I N Cl H O F” Chemical Bonds Compound – substance composed of two or more elements which are chemically combined in a fixed, specific ratio by mass Two general ways atoms chemically combine: 1. Via transfer of electrons 2. Via sharing of electrons Ionic Occurs between metal and nonmetal Metal loses electrons to nonmetal - Positively charged cations (from the metal) - Negatively charged anions (from the nonmetal) These ions bind to one another via electrostatic attraction – “opposites attract” Crystal Lattice – an array of attending ions/atoms Coulomb’s Law Attraction/repulsion increases with… - Decreasing ion size - Increasing ion charge Force (directionally proportional) (charge 1)(charge 2) / distance - - Metals – tend to lose e until they have the same number of e as the nearest noble gas Ex: Na (like Ne), Ca (like Ar) Nonmetals – tend to gain e until they have the same number of e as the nearest noble gas Covalent Bonding Occurs between nonmetal atoms Sharing electrons between them NO ions present in bond Forms when “electron clouds” of adjacent atoms overlap Chemical Formulas Consists of: Chemical symbols Numerical subscripts Indicates: Type of each atom/ion present (via chemical symbols) Number of each atom/ion present (via numerical subscripts) In the smallest unit of a substance Nomenclature Name compounds based on what type of bonds are predominantly present in the compound: Ionic compounds (combination of ions) Covalent compounds (nonmetal + nonmetal) Binary Covalent Compounds Prefix – “nonmetal #2” Prefix – “nonmetal #2” – ide Chemical formula represents smallest unit, the molecule. Prefixes indicate the number of atoms of each element present in the molecule. Number Prefix 1 Mono- 2 Di- 3 Tri- 4 Tetra- 5 Penta- 6 Hexa- 7 Hepta- 8 Octa- 9 Nona- 10 Deca- Don’t use “mono” with the FIRST element in a covalent compound. Ex: N2 3– dinitrogen trioxide, sulfur hexafluoride6– SF Binary Ionic Compounds “cation name” “anion name” Chemical formula represents the formula unit Represents the smallest, whole number ratio of ions that result in a net charge of ZERO NOT MOLECULES Prefixes are NOT used, as the charges of the ions indicate the relative ratio of elements (ions). Group 1A: 1+ Nonmetals Group 2A: 2+ Group # - 8 Aluminum: 3+ Group 7A: 1- Zinc: 2+ Group 6A: 2- Silver: 1+ Group 5A: 3- Cross charges – reduce to smallest numbers Cations with more than one charge – must indicate the compound name using Roman numerals Ex: “iron oxide” – Fe O (Iron (III) Oxide) or FeO (Iron (II) Oxide) 2 3 Polyatomic Ions “cation name” “anion name” Some molecules have an overall charge – polyatomic ions Can be cations or anions Since they are molecules, they stay covalently bonded Consists of 2 or more atoms covalently bonded together and has an overall charge Memorize: Ammonium (NH ) 4+ Hydroxide (OH ) - - Chlorate (ClO 3 - Nitrate (NO 3 Carbonate (CO ) 32- Phosphate (PO ) 43- Sulfate (SO 42- ONLY use parentheses if there are more than one polyatomic ion Acids H X (A = polyatomic or nonmetal) Form of covalent compound where hydrogen atoms ionize and become hydrogen cations H → H + e - Binary Acids Hydro – “nonmetal” – ic acid Gaseous compounds that become acids when dissolved in water (hence “hydro”) Symbol (aq) indicates “in a solution of water” Ex: HCl (aq) → hydrochloric acid, hydrosulfuric acid → H S 2aq) Ternary Acids “polyatomic” – ic acid No hydro prefix No (aq) required Do not require aqueous environment to be considered acids Also referred to as “oxy-acids” or “oxoacids” Hydrates M X Y nH O 2 “cation name” “anion name” prefix – hydrate Hydrates – compounds in solid state whose lattices contain one or more molecules per molecule/formula unit Anhydrous – compounds devoid of such water “impurities” Ex: Ba(OH) 2 8H O 2 barium hydroxide octahydrate, calcium sulfate dehydrate – CaSO ● 2H O 4 2
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