Chem 115 test 1
Symbols of Elements
Elements are symbolized by one or two letters; first letter is capitalized
Sig Fig Rules for Multiplication and Division
When multiplying or dividing numbers, the answer reported can not have more significant figures than either of the original numbers.
Periodicity
When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities; properties periodically repeat so that atoms in a group will have similar properties
Allotropes of Phosphorous
White phosphorus - crystalline solid P4 Red phosphorus - polymeric solid
Allotropes of Oxygen
dioxygen O2 - colorless ozone O3 - blue(absorbs UV rays, lighter than air so it sits in the atmosphere)
Calculating Empirical Formulas (look at notes for examples)
mass % elements --->grams of each element(assume 100g sample)--->moles of each element(use molar mass)--->calculate mole ratio to find the empirical formula
prefixes for binary compounds
mono- 1 di- 2 tri- 3 tetra- 4 penta- 5 hexa- 6 hepta- 7 octa- 8 nona- 9 deca- 10
Avogadro's Number
• 6.02 x 10^23 •= the number of particles in 1 mole of any substance • 1 mole of 12C has a mass of 12 g
Mixtures
=combinations of 2 or more substances in which each substances retains its chemical identity -the compositions of a mixture can vary -the substances making up a mixture are called components
Calculating the molecular formula from the empirical formula
*the subscripts in th emolecular formula of a a substance are always whole number multiples of the subscripts in its empirical formula Whole-number multiple= molecular weight/empirical formula weight (when given the MW) when we find this whole-number multiple we multiply the subscripts in the empirical formula by that number and we get the molecular formula
Cathode Ray tube exp
- used to study electrical discharge -a glass tube was pumped almost empty of air and a high voltage was applied to the electrodes in the tube, radiation was produced b/w the electrodes (radiation called cathode rays) -the radiation originated at the negative electrode(cathode) to the positive electrode(anode) -the presence of the rays were detected b/c they cause certain materials to fluoresce -Experiments showed that cathode rays are deflected by electric or magnetic fields.. proving that the cathode rays are streams of negatively charged particles -JJ Thomson observed that cathode rays are the same regardless of the identity of the cathode material; his paper is generally accepted as the discovery of the electron -Using a cathode-ray tube JJ Thomson determined that ratio of the electron's electrical charge to its mass
Atomic Number
-All atoms of the same element have the same number of protons -Atomic number=the number of protons not the number of electrons b/c electrons are used in reactions so the number of electrons is always changing, number of protons do not change -The atomic number (Z), it is the subscript
Octet rule
-Atoms want 8 electrons in their outer shell to be stable(or 2 electrons if the first ring is their outermost ring) -group 1A, 2A, and 3A will give up electrons -groups 5A, 6A, and 7A will gain electrons
Ionic Bonds
-Ionic compounds (such as NaCl) are generally formed between metals and nonmetals. -to form an ionic bond, one atom(metal) will lose an electron and the other atom(nonmetal) will gain the electron(these bonds form due to a give an take of electrons) -Ionic bonds are an electrostatic attraction -When an atom loses an electron the atom becomes smaller in size; When an atoms gains an electron is becomes larger in size -All ionic compounds are crystalline solids with high MP (MP above room temp)
Measuring Atomic Mass
Atomic and molecular masses can be measured with great accuracy with a mass spectrometer.
Law of Multiple Proportions
-It states that when elements combine they do so in a ratio of small whole numbers. For example, carbon and oxygen react to form CO or CO2, but not CO1.8. -When two elements combine with each other to form two or more compounds, the ratios of the masses of one element that combines with the fixed mass of the other are simple whole numbers.... ex: Carbon and Oxygen can combine to form CO or CO2 , the oxygen compared b/w these 2 compounds has a simple 1:2 ratio; a set 100 g of Carbon in both compounds will react with either 133g of Oxygen to produce CO or with 266g of Oxygen to produce CO2 -It is one of the basic laws of stoichiometry -Part of Dalton's atomic theory
Chromatography
-Separates substances on the basis of differences in solubility in a solvent; each substance dotted on paper will separate and move up paper depending on their solubility... the more soluble the substance the higher up the paper the substance moves - There is gas, liquid, and paper chromatography
Law of Constant Composition(aka the Law of Definite Proportions)
-The elemental composition of a pure substance never varies. -All samples of a given chemical compound have the same elemental composition by mass. For example, oxygen makes up about 8/9 of the mass of any sample of pure water, while hydrogen makes up the remaining 1/9 of the mass.
nonmetals
-The elements in groups 4A (14) through 7A (17) that lie above the stair step on some periodic tables. -Individual nonmetals may be either solids, liquids, or gases at room temperature. - They are poor conductors of heat and electricity. -Nonmetals characteristically are anions, when existing as monatomic ions in ionic compounds. When combined with other nonmetals, they typically form molecular compounds or complex ions.
Atomic Mass
-The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom. -It is the superscript
metals
-elements in groups 1A (1) and 2A (2), the transition elements, the lanthanides and actinides, and the heavier elements in groups 3A (13) through 5A (15) that lie below the stair step shown on some periodic tables. -At room temperature metals are shiny solids (except mercury and gallium above 29.78oC, which are liquids) that are malleable, ductile, and conductive of heat and electricity. -Metals characteristically are cations in their ionic compounds.
Common Anions
-monoatomic ions end in "ide"
Common Cations
-most transition metals form more than one type of cation with the EXCEPTION of Zinc(Zn2+) and Silver(Ag+)
metalloids
-the elements B, Si, Ge, As, Sb, Te, Po, At, which lie along the stair step shown on some periodic tables. -All are solids with semi-metallic properties. -They show poorer conductivity relative to metals and may be semiconductors (e.g., Si and Ge).
Ch 2: Atoms, Molecules, and Ions
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Ch 3: Stoichiometry: Calculations with Chemical Formulas and Equations
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5 Groups
1) Alkali Metals: Li, Na, K, Rb, Cs, Fr 2) Alkaline Earth Metals: Be, Mg, Ca, Sr, Ba, Ra 3) Chalcogens: O, S, Se, Te, Po 4) Halogens: F, Cl, Br, I, At 5) Noble Gases: He, Ne, Ar, Kr, Xe, Rn
Dalton's Postulates
1) All matter is composed of atoms. 2) All atoms of an element have the same mass (atomic weight). 3) All atoms of different elements have different masses (i.e., different atomic weights). 4) Atoms are indestructible and indivisible(Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.) 5) Compounds are formed when atoms of two or more elements combine. 6) In a compound the relative numbers and kinds of atoms are constant. #2,3, and 4 are now known to be incorrect because: 2) Many elements are composed of a mixture of isotopes, atoms of the same element with different masses. 3) Some atoms of two different elements may have virtually the same mass; these are called isobars. 4) Atoms can be split (fission) or merged (fusion) in nuclear reactions. Some of the mass of atoms is converted to energy in nuclear reactions; they are not indivisible because they are made up of protons, electrons, and neutrons
states of matter
1) Gas: has no fixed volume or shape, it conforms to the volume and shape of its container; the molecules are far apart and move at high speeds , colliding repeatedly with each other and with the walls of the container 2)Liquid: has a distinct volume independent of its container but has no specific shape; molecules are packed closely together but still move rapidly, and the molecules slide over one another 3)Solid: has both a definite shape and volume; molecules are held tightly together, usually in a definite arrangement in which the molecules can slightly wiggle *neither solids nor liquids can be compressed to any appreciable extent
When the first element in a compound is a nonmetal the second element in the compound will be a ? and will have a ? bond except for ?
1) nonmetal 2)covalent 3)NH4OH; NH4Cl (have ionic bonds)
Three types of radiation are? ;discovered by?
1) α particles(fast moving particles, more positively charged and heavier) 2) β particles(fast moving particles, more negatively charged and lighter: shown by the fact tat is deflects more) 3) γ rays(neutral; has the most energy(high-energy radiation) and will travel further) discovered by: Ernest Rutherford
4 common acids to know
1)Acetic acid, CH3COOH(organic acid) 2)Phosphoric acid, H3PO4 3)Chlorous acid, HClO2 4)Hypochlorous acid, HClO
Isotopes of Hydrogen
1H: normal hydrogen with one proton and no neutrons; called Protium 2H: Hydrogen with one proton and 1 neutron; called Deuterium(Heavy water) 3H: Hydrogen with 1 proton and 2 neutrons; called Tritium -All three of these have identical chemical properties; chemical properties depend on electrons
Molecules
=2 or more atoms joined together in a specific shape -Molecules of elements are homonuclear, because they are composed of only one kind of atom. -Molecules of compounds are heteronuclear, because they are composed of two or more different kinds of atoms.
Temperature
=A measure of the average kinetic energy of the particles in a sample; it is a physical property • In scientific measurements, the Celsius and Kelvin scales are most often used. • The Celsius scale is based on the properties of water. □ 0°C is the freezing point of water. □ 100°C is the boiling point of water. • The Kelvin is the SI unit of temperature; It is based on the properties of gases; There are no negative Kelvin temperatures, it starts at absolute zero • K=°C+273.15
Matter
=Anything that has mass and takes up space • Atoms are the building blocks of matter.(atoms= smallest unit of matter) • Each element is made of the same kind of atom. • A compound is made of two or more different kinds of elements
nuclide
=any atom with a certain number of nucleons(protons+neutrons)
Property
=any characteristic that allows us to recognize a particular type of matter and to distinguish it from other types
Compounds
Compounds can be broken down into more elemental particles; a compound will have different properties than the individual elements on their own
network solids
=elements or compounds in which all the atoms are bound together in a limitless 3D structure (ex: the 2 forms of carbon, graphite and diamond); b/c there are no discrete molecules, the formua for a network solid can only be represented as an empirical formula
Solutions
=homogenous mixtures; solutions can be solids, liquids, or gases
Density
=m/V -density is temp dependent b/c most substances change vol when heated or cooled
Pure substance
=matter that has distinct properties and a composition that does not vary from sample to sample
Polyatomic Ions
=molecular ions with specific charges •Ions that contain more than 1 atom; can be anions or cations(there are only 2 polyatomic cations= NH4+(ammonium) and Hg2 2+... the rest are anions) •Ex: NH4+, CO3 2- •within the polyatomic ion the elements are covalently linked but they still hav an overall charge and when they form a compound with another ion it forms an ionic bond
Isobars
=nuclides of different elements, Different atomic number with same atomic mass
Elements
=substances that cannot be broken down into simpler substances; each element is composed of only one type of atom
Scientific Method
=systematic approach to solving problems -observation and experiments; find patterns, trends, and laws; formulate and test hypothesis; develop a theory
7 strong acids
HCl= hydrochloric HBr= hydrobromic HI= hydroiodic HClO3= chloric HClO4= perchloric HNO3= nitric H2SO4= sulfuric (both hydrogens are acidic) *all acids are covalent compounds but will dissociate in water
Allotropes
Many elements like C, O and S exist in more than one form in nature and these are called allotropes
Percent Composition
One can find the percentage of the mass of a compound that comes from each of the elements in the compound by using this equation: % element =(number of atoms * atomic weight)/ (FW of the compound) x 100
What is the most abundant element on earth?
Oxygen
Metric System
Prefixes convert the base units into units that are appropriate for the item being measured. •Giga(G), 10^9 •Mega(M), 10^6 •Kilo(k), 10^3 •Deci(d), 10^-1 •Centi(c), 10^-2 •Milli(m), 10^-3 •Micro(mu), 10^-6 •Nano(n), 10^-9 •Pico(p), 10^-12 •Femto(f), 10^-15
Distillation
Separates homogeneous mixture on the basis of differences in boiling point.
Filtration
Separates solid substances from liquids and solutions.
Chemistry
The study of matter and the changes it undergoes.
Atomic Theory of Matter
The theory that atoms are the fundamental building blocks of matter; reemerged in the early 19th century, championed by John Dalton.
Law of Conservation of Mass
The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place
Sig Fig Rules for Addition and Subtraction
When adding or subtracting numbers, the reported answer can not have more digits after the decimal point than any of the added numbers.
Periodic Table
• A systematic catalog of elements. • Elements are arranged in order of atomic number. •Based on electron configuration • The rows on the periodic chart are periods. • Columns are groups; the ones labeled with an A (1A-8A) are the main groups and the number represents the number of valence electrons in the outer shell; IUPAC order of numbering groups is 1-18 • Elements in the same group have similar chemical properties. • Mendeleev's table is considered the first periodic table, published in 1869 with many gaps and uncertainties and he made predictions that later on more elements would be discovered that will fill in the gaps
Accuracy versus Precision
• Accuracy refers to the proximity of a measurement to the true value of a quantity. • Precision refers to the proximity of several measurements to each other (the repeatability of measurements); precision of measurement is is often expressed in terms of the standard deviation
Molar Mass
• By definition, these are the mass of 1 mol of a substance (i.e., g/mol) - The molar mass of an element is the mass number for the element that we find on the periodic table - The formula weight (in amu's) will be the same number as the molar mass (in g/mol)
basic history need to know:
• Dalton: Atomic theory • J J Thomson: Electrons and their charge to mass ratio • Millikan: Charge on the electron and later on the mass • Henry Becquerel: Radioactivity • Rutherford: α β and γ radiation; Gold leaf experiment to show the presence of nucleus and coined the word the proton • James Chadwick: Neutrons
Allotropes of Carbon
• Diamond - an extremely hard, transparent crystal, with the carbon atoms arranged in a tetrahedral lattice. A poor electrical conductor. An excellent thermal conductor. • Graphite - a soft, black, flaky solid, a moderate electrical conductor. The C atoms are bonded in flat hexagonal lattices, which are then layered in sheets.(found in pencils) • amorphous carbon • fullerenes including "buckyballs"(spherical), such as C60, and carbon nanotubes
Atomic Structure
• Discovery of Electrons: J J Thomson ( in 1897) • He gave the charge to mass ratio of electrons. • He also showed that the nature of the cathode ray did not depend on the material of the cathode. • Conclusions: 1. Cathode ray consists of negatively charged particles and electrons are constituents of all matter. 2. The charge on electrons is found to be 1.602 x 10 -19 coulombs.
Empirical vs Molecular formulas
• Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound; ex: CH2O (empirical formula for glucose); all ionic compounds are represented by their empirical formula, ex: NaCl • Molecular formulas give the exact number of atoms of each element in a compound; ex: C6H12O2 (glucose)
Acid Nomenclature
• If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro- Ex: HCl: hydrochloric acid HBr: hydrobromic acid HI: hydroiodic acid • If the anion in the acid ends in -ate, change the ending to -ic acid Ex: HClO3: chloric acid HClO4: perchloric acid • If the anion in the acid ends in -ite, change the ending to -ous acid: Ex: HClO: hypochlorous acid HClO2: chlorous acid •Anions derived by adding H+ to the oxyanion you just add the word hydrogen at the beginning; Ex: HCO3-, hydrogen carbonate
Classification of matter
• If the matter is uniform throughout then it is homogeneous, if it is not uniform throughout then it is a heterogenous mixture • Homogeneous: 1)if it has a variable composition(meaning it can be separated by physical means) then it is called a homogeneous mixture (solution); if it doesn't have a variable composition then it is called a Pure Substance 2)if the pure substance can be separated into simpler substances(by chemical means) then it is a compound; if the pure substance cannot be separated into simpler substances then it is an element
Nomenclature of Binary Compounds
• Naming for a compound containing 2 nonmetals • The less electronegative atom is usually listed first.(exception: when the compund contains oxygen, and chlorine, Bromine, or Iodine then the oxygen is written last) • A prefix is used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however.) • The ending on the more electronegative element is changed to -ide. • If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one: N2O5= dinitrogen pentoxide Ex: CO2: carbon dioxide CCl4: carbon tetrachloride •Exception: molecular compuns that cotain H and one other element can be treated as if they were neutral substances containing a H+ ion and anion... thus you can predict that the substance named hydrogen chloride is HCl, containing one H+ and one Cl-(the name hydrogen chloride is only used for the pure compund, water solns of HCl are called hydrochloric acid) Similarily H2S is called Hydrogen Sulfide
Discovery of neutrons
• Neutrons were discovered by James Chadwick in 1932.
Discovery of the location of protons (the "nuclear model" of an atom)
• Now since a negative particle was discovered it was definite that the atom would also have an equal positive charge. • Then at the time the most reasonable explanation seemed to be a plum pudding model(aka a random distribution of electrons and protons • Rutherford explained the existence of a focused positively charged nucleus in the atom; Protons were discovered by Rutherford in 1919. • the experiment he performed was he bombarded a gold thin foil with alpha particles(used gold b/c it can be hammered into very thin sheets); the result was that most of the alpha particles went straight thru the gold foil, some were deflected(alpha particles that were close to the nucleus were deflected )... conclusion: positive charge is focused in a specific area of the atom not randomly dispersed • Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom. • Most of the volume of the atom is empty space.
Millikan Oil Drop Experiment
• Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would yield the other. • Robert Millikan (University of Chicago) determined the charge on the electron in 1909. •Using the charge and Thomson's charge to mass ratio of an electron, Millikan calculated the mass of an electron
Decomposition Reactions (REACTION 2)
• One substance breaks down into two or more substances • Examples: CaCO3 (s) → CaO (s) + CO2 (g) 2 KClO3 (s) → 2 KCl(s) +3 O2 (g) 2NaN3(s) → 2Na(s) +3N2(g)
Changes of Matter
• Physical Changes= Changes in matter that do not change the composition of a substance, can change its appearance Ex: Changes of state, temperature, volume, etc. • Chemical Changes= Changes that result in new substances. Ex: Combustion, oxidation, decomposition, etc.
Properties of Matter
• Physical Properties= Can be observed without changing a substance into another substance(without changing the identity and composition of the substance) Ex: Boiling point, density, mass, volume, color, odor, melting point, hardness, etc. • Chemical Properties=Can only be observed when a substance is changed into another substance. Ex: Flammability, corrosiveness, reactivity with acid, etc. • Intensive Properties= Independent of the amount of the substance that is present. Ex: Density, boiling point, color, etc. • Extensive Properties= Dependent upon the amount of the substance present. Ex: Mass, volume, energy, etc.
Allotropes of Sulfur
• Plastic(amorphous)sulfur: polymeric solid • Rhombic sulfur: large crystals composed of S8 molecules • Monoclinic sulfur: fine needle-like crystals • Other ring molecules such as S7 and S12
Subatomic Particles
• Protons and electrons are the only particles that have a charge. • Protons and neutrons have essentially the same mass(1 amu) •The mass of an electron is so small we ignore it.
Combustion Reactions (REACTION 3)
• Rapid reactions that produce a flame • Most often involve hydrocarbons reacting with oxygen in the air • Examples: CH4(g) +2O2(g) → CO2(g) +2H2O(g) C3H8(g)+ 5 O2(g) → 3 CO2(g) + 4 H2O(g)
Anatomy of a Chemical Equation
• Reactants appear on the left side of the equation. • Products appear on the right side of the equation. • The states of the reactants and products are written in parentheses to the right of each compound. (g)(l)(s)(aq) (aq)= compound dissolved in water • Coefficients are inserted to balance the equation.
Structural formulas vs Perspective drawings
• Structural formulas show the order in which atoms are bonded; flat • Perspective drawings also show the three-dimensional array of atoms in a compound; a wedge points out of the page, a dotted line points into the page, straight lines are on the plane of the page •There are also the Ball-and Stick model, and the Space-filling model
Formula Weight
• Sum of the atomic weights for the atoms in a chemical formula • So, the formula weight of calcium chloride, CaCl2, would be... Ca: 1(40.1 amu) + Cl: 2(35.5 amu)=111.1 amu • These are generally reported for ionic compounds which are represented by their empirical formulas
Molecular Weight (MW)
• Sum of the atomic weights of the atoms in a molecule(used for covalent compounds) • For the molecule ethane, C2H6, the molecular weight would be... C: 2(12.0 amu) + H: 6(1.0 amu)= 30.0 amu
Volume
• The most commonly used metric units for volume are the liter (L) and the milliliter (mL).; volume is a derived SI unit, it is derived from the SI unit for length=m, volume=m^3 □ A liter is a cube 1dm long on each side. □ A milliliter is a cube 1 cm long on each side.
Radioactivity
• The spontaneous emission of radiation by an atom. • An atom is radioactive when the nucleus is unstable and so it is emitting particles from the nucleus • All radioactive elements will decay into lead , the nucleus of lead is more stable •Uranium is the heaviest unstable, naturally occurring atom, it is radioactive; Bismuth is the heaviest stable, naturally z atom(after Bi the neutrons cannot accommodate for the increase in protons and the atoms become radioactive) • First observed by Henri Becquerel. • Also studied by Marie and Pierre Curie.
Significant Figures
• The term significant figures refers to digits that were measured. • When rounding calculated numbers, we pay attention to significant figures so we do not overstate the accuracy of our answers. • Rules: 1.All nonzero digits are significant. 2.Zeroes between two significant figures are themselves significant. 3.Zeroes at the beginning of a number are never significant. 4.Zeroes at the end of a number are significant if a decimal point is written in the number. *For numbers written in scientific notation, All digits in the coefficient are significant
Average Mass
• The unit for atomic mass is amu • Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations. • Average mass(aka atomic weight) is calculated from the isotopes of an element weighted by their relative abundances.
Combination Reactions (REACTION 1)
• Two or more substances react to form one product • Examples: N2(g) +3H2(g) → 2NH3(g) C3H6 (g) + Br2 (l) → C3H6Br2 (l) 2Mg(s) +O2 (g) → 2MgO(s) this is a metal reacting with a nonmental to prodice an ionic solid
Patterns in Oxyanion Nomenclature
• When there are two oxyanions involving the same element: -The one with fewer oxygens ends in -ite (Ex: NO2 − : nitrite; SO3 2− : sulfite; ClO2-: chlorite ) -The one with more oxygens ends in -ate (Ex: NO3 − : nitrate; SO4 2− : sulfate; ClO3-: chlorate) •When a compound has more than 2 oxyanions then you use the prefixes hypo and per -The one with the fewest oxygens has the prefix hypo- and ends in -ite; Ex: ClO− : hypochlorite -The one with the most oxygens has the prefix per- and ends in -ate; ex: ClO4− : perchlorate; the prefix "per-" indicates that a chemical compound contains a high proportion of a specific element •A way to remember some polytomic ions: C and N, both period 2 elements, have only 3 O atoms whereas period 3 elements, P, S, and Cl, have 4 O atoms each; among these atoms, ionic charge increases from right to left(ex: 1- for ClO4^- to 3- for PO4^3-; NO3^- to CO3^2-)
Inorganic nomenclature for a cation and anion
• Write the name of the cation then the word "ion" after it ; If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses.(ex: Iron(II) ion, Iron(III) ion) • Older method for naming charged ions of metals uses the endings -ous and -ic added to the root of the latin name (ex: Ferrous ion (Fe2+) and Ferric ion(Fe3+)) • If the anion is an element(monoatomic), change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion.
Isotopes
•Atoms of the same element with different masses. • Isotopes have different numbers of neutrons but the number of protons is the same • Every atom except F has isotopes; isotopes are radioactive
Writing Formulas
•Because compounds are electrically neutral, one can determine the formula of a compound this way: -The charge on the cation becomes the subscript on the anion. -The charge on the anion becomes the subscript on the cation. -If these subscripts are not in the lowest whole-number ratio, divide them by the greatest common factor.
Strong Bases
•Group 1A metals(Alkali metals) hydroxides: LiOH, NaOH, KOH, RbOH, CsOH •Heavy group 2A metal hydroxides: Ca(OH)2, Sr(OH)2, Ba(OH)2
Diatomic Molecules
•Include H2, N2, O2, F2, Cl2, Br2, and I2 •These seven elements occur naturally as molecules containing two atoms; the pure is substance is found as the diatomic molecule(ex: pure substance of hydrogen is H2 not H)
SI Units (7 base units)
•Mass, kg •Length, m •Time, s •Temp, K •Amount of a substance, mol •Electric current, Ampere(A) •Luminous intensity, Candela(cd)
Molecular Compounds
•Molecular compounds are composed of molecules and almost always contain only nonmetals. •Metals will not form compounds with other metals they will just form a mixture, called an alloy •Metals can form compounds with nonmetals called ionic compounds , not a molecule b/c there is not a covalent bond b/w the atoms •Nonmetals can form compounds with nonmetals called covalent compounds, this is a true molecule
Periodic Table(metals, nonmetals, metalloids)
•Nonmetals are on the right side of the periodic table (with the exception of H). •Metalloids border the stair-step line (with the exception of Al and Po); include B, Si, Ge, As, Sb, Te, At... have properties of both metals and nonmetals •Metals are on the left side of the chart; most are found as compounds in nature in order to be stable, by themselves they are unstable(aka highly reactive)... EXCEPTION: Gold(Au) and noble gases are found on their own in nature
Mole Relationships
•One mole of atoms, ions, or molecules contains Avogadro's number of those particles • One mole of molecules or formula units contains Avogadro's number times the number of atoms or ions of each element in the compound
Subscripts vs Coefficients
•Subscripts tell the number of atoms of each element in a molecule • Coefficients tell the number of molecules
Exact numbers
•These numbers are the ones whose values are known exactly which include all integer fractions(1/2, 1/3, 7/8), Counted numbers, and Conversion factors(conversions within a unit system!) ***Relationships b/w units in different unit systems are usually NOT exact, EX: 2.2lb=1kg has 2 sig figs... EXCEPTION: 2.54cm= 1in (exact), 1cal= 4.184J (exact) •The numbers that are obtained by counting and not by measuring are called exact numbers. • Exact numbers can be assumed to have an unlimited number of significant figures; These do not limit the number of significant figures in a calculation.
Ions
•When atoms lose or gain electrons,they become ions. -Cations are positive and are formed by elements on the left side of the periodic chart(metals)(group 1A, 2A, and 3A will give up electrons) -Anions are negative and are formed by elements on the right side of the periodic chart(non-metals) (groups 5A, 6A, and 7A will gain electrons) -Ionic compounds are typically crystalline solids with a structure consiting of an orederly 3D array of ions called a crystal lattice *Group 4A cannot gain 4 or lose 4 electrons therefore they will form covalent bonds instead of ionic bonds
