AP Ch 2 Atoms Molecules Ions Nomenclature

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Atoms, Molecules and Ions: 

Atoms, Molecules and Ions Chapter 2

Dalton’s Atomic Theory (1808): 

Dalton’s Atomic Theory (1808) Elements are composed of extremely small particles called atoms. All atoms of a given element are identical, having the same size, mass and chemical properties. The atoms of one element are different from the atoms of all other elements. Compounds are composed of atoms of more than one element. The relative number of atoms of each element in a given compound is always the same. Chemical reactions only involve the rearrangement of atoms. Atoms are not created or destroyed in chemical reactions. 2.1


2 2.1


8 X2Y 2.1


J.J. Thomson, measured mass/charge of e- (1906 Nobel Prize in Physics) 2.2


Cathode Ray Tube 2.2


e- charge = -1.60 x 10-19 C Thomson’s charge/mass of e- = -1.76 x 108 C/g e- mass = 9.10 x 10-28 g Measured mass of e- (1923 Nobel Prize in Physics) 2.2

Everybody Has Avogadro’s Number! But Where Did it Come From?: 

Everybody Has Avogadro’s Number! But Where Did it Come From? It was NOT just picked! It was MEASURED. One of the better methods of measuring this number was the Millikan Oil Drop Experiment Since then we have found even better ways of measuring using x-ray technology


(Uranium compound) 2.2




The modern view of the atom was developed by Ernest Rutherford (1871-1937).


atoms positive charge is concentrated in the nucleus proton (p) has opposite (+) charge of electron (-) mass of p is 1840 x mass of e- (1.67 x 10-24 g) particle velocity ~ 1.4 x 107 m/s (~5% speed of light) (1908 Nobel Prize in Chemistry) 2.2


atomic radius ~ 100 pm = 1 x 10-10 m nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m Rutherford’s Model of the Atom 2.2 “If the atom is the Houston Astrodome, then the nucleus is a marble on the 50-yard line.”

Chadwick’s Experiment (1932): 

Chadwick’s Experiment (1932) H atoms - 1 p; He atoms - 2 p mass He/mass H should = 2 measured mass He/mass H = 4 neutron (n) is neutral (charge = 0) n mass ~ p mass = 1.67 x 10-24 g 2.2


mass p = mass n = 1840 x mass e- 2.2


Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei 2.3




6 protons, 8 (14 - 6) neutrons, 6 electrons 6 protons, 5 (11 - 6) neutrons, 6 electrons Do You Understand Isotopes? 2.3




Chemistry In Action Natural abundance of elements in Earth’s crust Natural abundance of elements in human body 2.4


A molecule is an aggregate of two or more atoms in a definite arrangement held together by chemical bonds A diatomic molecule contains only two atoms H2, N2, O2, Br2, HCl, CO A polyatomic molecule contains more than two atoms O3, H2O, NH3, CH4 2.5


ELEMENTS THAT EXIST AS DIATOMIC MOLECULES Remember: BrINClHOF These elements only exist as PAIRS. Note that when they combine to make compounds, they are no longer elements so they are no longer in pairs!


An ion is an atom, or group of atoms, that has a net positive or negative charge. cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. 2.5

Forming Cations & Anions: 

Forming Cations & Anions A CATION forms when an atom loses one or more electrons. An ANION forms when an atom gains one or more electrons Mg --> Mg2+ + 2 e- F + e- --> F-


A monatomic ion contains only one atom A polyatomic ion contains more than one atom 2.5 Na+, Cl-, Ca2+, O2-, Al3+, N3- OH-, CN-, NH4+, NO3-


13 protons, 10 (13 – 3) electrons 34 protons, 36 (34 + 2) electrons Do You Understand Ions? 2.5






A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance An empirical formula shows the simplest whole-number ratio of the atoms in a substance H2O C6H12O6 CH2O O3 O N2H4 NH2 2.6


ionic compounds consist of a combination of cation(s) and an anion(s) the formula is always the same as the empirical formula the sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero The ionic compound NaCl 2.6


Formula of Ionic Compounds Al2O3 2.6 Al3+ O2- CaBr2 Ca2+ Br- Na2CO3 Na+ CO32-





Examples of Older Names of Cations formed from Transition Metals (memorize these!!): 

Examples of Older Names of Cations formed from Transition Metals (memorize these!!) From Zumdahl

Chemical Nomenclature: 

Chemical Nomenclature Ionic Compounds often a metal + nonmetal anion (nonmetal), add “ide” to element name BaCl2 barium chloride K2O potassium oxide Mg(OH)2 magnesium hydroxide KNO3 potassium nitrate 2.7


Transition metal ionic compounds indicate charge on metal with Roman numerals FeCl2 2 Cl- -2 so Fe is +2 iron(II) chloride FeCl3 3 Cl- -3 so Fe is +3 iron(III) chloride Cr2S3 3 S-2 -6 so Cr is +3 (6/2) chromium(III) sulfide 2.7


Molecular compounds nonmetals or nonmetals + metalloids common names H2O, NH3, CH4, C60 element further left in periodic table is 1st element closest to bottom of group is 1st if more than one compound can be formed from the same elements, use prefixes to indicate number of each kind of atom last element ends in ide 2.7


HI hydrogen iodide NF3 nitrogen trifluoride SO2 sulfur dioxide N2Cl4 dinitrogen tetrachloride NO2 nitrogen dioxide N2O dinitrogen monoxide Molecular Compounds 2.7




An acid can be defined as a substance that yields hydrogen ions (H+) when dissolved in water. HCl Pure substance, hydrogen chloride Dissolved in water (H+ Cl-), hydrochloric acid An oxoacid is an acid that contains hydrogen, oxygen, and another element. 2.7 HNO3








A base can be defined as a substance that yields hydroxide ions (OH-) when dissolved in water. 2.7



Mixed Practice: 

Mixed Practice Dinitrogen monoxide Potassium sulfide Copper (II) nitrate Dichlorine heptoxide Chromium (III) sulfate Ferric sulfite Calcium oxide Barium carbonate Iodine monochloride N2O K2S Cu(NO3)2 Cl2O7 Cr2(SO4)3 Fe2(SO3)3 CaO BaCO3 ICl

Mixed Practice: 

Mixed Practice BaI2 P4S3 Ca(OH)2 FeCO3 Na2Cr2O7 I2O5 Cu(ClO4)2 CS2 B2Cl4 Barium iodide Tetraphosphorus trisulfide Calcium hydroxide Iron (II) carbonate Sodium dichromate Diiodine pentoxide Cupric perchlorate Carbon disulfide Diboron tetrachloride

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