Chemistry - Ch 2 Flashcards
Law of constant composition
In a given compound, the relative numbers and kinds of atoms are constant
Atoms
Smallest particles of an element that retain the chemical identity of the element
Law of conservation of mass/matter
Total mass of materials present after a chemical reaction is the same as the total mass present b4 the reaction; basis for Postulate 3
Law of multiple proportions
If two elements A & B combine to form more than one compound, the masses of B can combine with a given mass of A are in the ratio of small whole numbers
Subatomic particles
Smaller parts of an atom
Cathode rays
Radiation originating from the negative electrode (cathode)
Coulomb
SI unit for electrical charge
Charge of a single electron
1.602 x 10^-19
Thomson’s charge to mass ratio
1.76 x 10^8 C/g
Mass of electron
9.10938 x 10^-28g; 2000 times smaller than hydrogen (lightest atom)
Radioactivity
Spontaneous emission of radiation; discovered by Henri Becquerel, who urged Marie Curie & her husband Pierre to isolate the radioactive components of uranium
3 types of radiation
Ernest Rutherford: alpha , beta, gamma (y); each type responds differently in an electric field - a & b are bent by it; y is unaffected
Beta rays
High speed electrons; = radioactive element of cathode rays; attracted to a positively charged plate
Alpha rays
Positive charge; attracted toward a negative plate
Gamma rays
High energy radiation similar to x-rays; does not consist of particles, carries no charge
Nucleus
Mass of atom & all of its positive charge reside in this very small dense region per Ernest Rutherford in his scattering experiment, which disproved JJ Thomson’s plum pudding model; protons & neutrons reside in nucleus
Discovery of protons (+ particles)
1919, Ernest Rutherford
Discovery of neutrons (neutral)
1932, James Chadwick
Proton charge (electronic charge)
+1.602 x 10^-19 C
Every atom has ___________ number of electrons & protons so atoms have no net electrical charge
Equal
1 Atomic mass unit (amu)
1.66054 x 10^-24g; protons & neutrons’ masses nearly equal, both much greater than an electron; proton’s mass: 1.0073 amu; neutron: 1.0087 amu; electron: 5.486 x 10^-4 amu; 1g= 6.02214 x 10^23amu
Angstrom
Non SI unit of length to express atomic dimensions
Atomic number
Number of protons in the nucleus of an atom
Mass number
Element symbol with subscript & superscript to left; superscript = mass # (protons + neutrons); subscript = atomic # (# of protons or electrons)
Isotopes
Atoms with identical atomic numbers but different mass numbers(same # of protons, different # of neutrons)
Average atomic mass
Masses of its various isotopes & their relative abundances
Mass spectrometer
Most direct and accurate means for determining atomic & molecular weights
Mass spectrum
Graph of the intensity of the detector signal vs particle atomic mass
Periodic table
Most significant tool chemists use for organizing & remembering chemical facts; arrangement of elements in order of increasing atomic number, with elements having similar properties placed in vertical columns
Periods
Horizontal rows of periodic table
Groups
Vertical columns of periodic table; AB designations with Arabic or Roman numerals, or IUPAC (Internat’l union of pure chemistry) 1-18 with no letters; similar properties b/c same arrangement of electrons @ periphery of their atoms
Coinage metals
Copper (cu), silver (ag), & gold (au); pt of gp 1B
Alkali metals (1A)
Li, Na, K, Rb, Cs, Fr
Alkaline earth metals (2A)
Be, Mg, Ca, Sr, Ba, Ra
Chalcogens (6A)
O, S, Se, Te, Po
Halogens (7A)
F, Cl, Br, I, At
Noble/rare gases (8A)
He, Ne, Ar, Kr, Xe, Rn
Metallic elements (metals)
All elements on left side and in middle of periodic table except hydrogen; share properties - luster, high electrical/heat conductivity; solid @ rm temp (except mercury)
Nonmetallic elements (nonmetals)
Separated from metals on table by diagonal steplike line from boron (B) to astatine (At); includes hydrogen; @ rm temp, some gaseous, some solid, 1 liquid; differ from metals in appearance & other physical properties
Metalloids
Elements that lie along line between metals & nonmetals that have properties that fall btwn both (ex.antimony-Sb)
Nuclear reactions
Reaction btwn uranium & neutrons that creates plutonium
Transuranium elements
Elements beyond uranium (92) on periodic table; not found in nature; can only be synthesized via nuclear reactions
Molecule
Assembly of 2+ atoms tightly bound together
Chemical formula
Represents molecular form of an element
Diatomic molecule
Molecule made up of 2 atoms; hydrogen (H2), oxygen (O2), nitrogen (N2) & the halogens (F2, Cl2, Br2, I2) normally occur as diatomic molecules
Ozone
O3; sharp, pungent smell, toxic
Molecular compounds
Compounds composed of molecules (contain more than 1 type of atom)
Molecular formulas
Chemical formulas that indicate actual numbers and types of atoms in a molecule
Empirical formulas
Chemical formulas that only give the relative number of atoms of each type of molecule; subscripts = smallest possible whole # ratios
Structural formula
Shows which atoms are attached to which within the molecule; atoms represented by chemical symbols, lines represent bonds
Perspective drawing
Drawing a structural formula to show the angles at which atoms are joined together
Ball-and-stick model
Show atoms as spheres & bonds as sticks; accurately represents angles of attachment
Space-filling model
Depicts what the molecule would look like if the atoms were scaled up in size; angles hard to see
Ion
Charged particle formed by removing or adding electrons to a neutral atom; positively charged ion = cation; negatively charged ion = anion
Metal ions _______ electrons to form cations; nonmetals ________ electrons to form anions
Lose;gain
Polyatomic ions
Atoms joined as a molecule but with a net positive or negative charge (ex. NH4+, ammonium ion; or SO4 2-, sulfate ion)
Ionic compound
Compound with both positively & negatively charged ions; in general cations are metal ions & anions are nonmetal ions; ionic compounds are generally combos of metals & nonmetals ex. NaCl (molecular compounds generally = nonmetals only ex.H2O); only empirical formulas can be written for most ionic compounds
Chemical nomenclature
System used in naming substances (Latin nomen - name - & calare - to call)
Organic compounds
Carbon, usually in combo with hydrogen, oxygen, nitrogen, or sulfur; all other compounds are inorganic
Monatomic ions
Ions formed from a single atom
Transition metals
Metals that can form more than one cation; occur in middle block of elements on periodic table, from group 3B to 2B; 1A (Na+, K+, & Rb+), 2A (Mg2+, Ca2+, Sr2+, Ba2+), 3A (Al3+) & 2 transition metal ions Ag+ (1B) & Zn2+ (2B) form only 1cation
Cations formed from nonmetal ions have names that end in _________
-ium ex. NH+4 (ammonium ion); H3O+ (hydronium ion); both ex are polyatomic; vast majority of cations are monatomic metal ions
Monatomic anions (neg ions) are formed by replacing ending with ______
IDE ex H- (hydride); O2- (oxide); N3- (nitride); a few simple polyatomic anions have names ending in IDE (OH- = hydroxide; CN- = cyanide; O2 2- = peroxide)
Oxyanions
Polyatomic anions containing oxygen; end in -ate (most common oxyanion of an element) or -ite (oxyanion that has same charge but 1 atom fewer (ex. NO3- = nitrate; NO2- = nitrite)
Anions derived by adding H+ to an oxyanion are named by adding as a prefix the word ____ or _____
Hydrogen or dihydrogen (ex. CO3 2- : carbonate ion; HCO3- : hydrogen carbonate; PO4 3- : phosphate ion; H2PO4 - : dihydrogen phosphate ion; each H+ reduces the negative charge of the parent ion by 1; older method for naming is to use prefix bi- (HCO3- = bicarbonate)
Prefix hypo
One O atom fewer than oxyanion ending in ite ex. ClO2- = chlorite; ClO- = hypochlorite (one O atom fewer than chlorite)
Ionic compounds consist of the ____name followed by the ____ name
Cation; anion (ex. CaCl2 = calcium chloride)
Acid
substance whose molecules yield hydrogen (h+) ions when dissolved in water; composed of an anion connected to enough H+ ions to neutralize/balance the anion’s charge
Acids containing anions whose names end in -ate or -ite or named by changing -ate to ___ and -ite to ___ then adding the word acid
-ic; -ous (Ex. ClO4 - : perchloric –> HClO4 (perchloric acid); ClO3 - : chlorate –> HClO3 (chloric acid); ClO2- : chlorite –> HClO2 (chlorous acid); ClO- : hypochlorite –> HClO (hypochlorous acid)
Names & Formulas of Binary Molecular Compounds - 3 Rules
(1) Name of element farther to left in table usually written 1st (exception = Oxygen; always last except when combined with fluorine (2) if both elements are in the same group, the one having the higher atomic # is named 1st (3) name of 2nd element is given an -ide ending (4) Greek prefixes are used to indicate # of atoms of each element; mono never used with 1st element; prefix ends in a or o & the name of 2nd begins with a vowel, a or o is often dropped
Prefixes for Naming Binary Compounds (10)
Mono (1), Di (2), Tri (3), Tetra (4), Penta (5), Hexa (6), Hepta (7), Octa (8), Nona (9), Deca (10)
Organic Chemistry
study of compounds of carbon
Hydrocarbons
Compounds that contain only carbon and hydrogen
Alkanes
each carbon atom is bonded to four other atoms (3 simplest = methane, CH4 =1 carbon; ethane, C2H6 = 2 carbons, & propane, C3H8 = 3 carbons); 4C = butane; 8 = octane (C8H18)
Functional groups
specific groups of atoms
Alcohol
Obtained by replacing an H atom of an alkane with an -OH group; replace -ane ending with -ol
