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Chapter 2 Atoms, Molecules, Ions HW: 4 11 23 25 31 35 39 45 50 53 55 59 61 63 65 67 69 71 102.

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Presentation on theme: "Chapter 2 Atoms, Molecules, Ions HW: 4 11 23 25 31 35 39 45 50 53 55 59 61 63 65 67 69 71 102."— Presentation transcript:

1 Chapter 2 Atoms, Molecules, Ions HW: 4 11 23 25 31 35 39 45 50 53 55 59 61 63 65 67 69 71 102

2 2.1 – Atomic Theory Dalton’s Atomic Theory 1.Elements are made of extremely small particles called atoms 2.Atoms of an element are identical 3.Atoms are not created or destroyed 4.Compounds are combinations of atoms (1766-1844)

3 2.1 – Atomic Theory Law of Definite Proportions (ConstantComposition) -Joseph Proust -A compound always has the same proportion of its elements Law of Conservation of Mass (Matter) = Matter cannot be created or destroyed Law of Multiple Proportions -Two or more compounds with the same elements must have different proportions of the elements -Example: Water vs. Hydrogen peroxide

4 2.2 – Discovery of Atomic Structure Atom = Basic unit of an element that can enter into a chemical reaction Subatomic Particle = Particles that make up an atom (protons, neutrons, electrons are as small as we will study)

5 -Cathode (Ray) Tube / Crooke’s Tube / = Glass tube w/2 metal plates. Connected to high voltage source. Emits ray. 2.2 – Discovery of Atomic Structure Crooke – Determined that the ray was made of negative particles

6 2.2 – Discovery of Atomic Structure JJ Thomson = 1897 – Credited w/ finding electrons Determined charge to mass ratio to be -1.76  10 8 coulombs/g. Millikan = 1909 - Performed the Oil Drop Experiment. Found the charge of an electron (-1.60 x 10 -19 C). Could then calculate mass of the electron

7 2.2 – Discovery of Atomic Structure  = + charge = Helium nucleus = High mass  = - charge (high speed electrons) = Low mass  = No charge. High energy = No mass Radioactivity = Spontaneous emission of radiation

8 2.2 – Discovery of Atomic Structure 1900 - “Plum pudding” model, put forward by JJ Thompson. Positive sphere of matter with negative electrons imbedded in it.

9 2.2 – Discovery of Atomic Structure Protons = + charge particles in the nucleus Rutherford = 1919- Gold Foil Experiment -Most alpha particles when through, some deflected

10 2.2 – Discovery of Atomic Structure This proved that Thomson’s model was incorrect Results of Gold Foil Experiment: 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.

11 2.2 – Discovery of the Structure of the Atom Mass of electron so small it is ignored -We will only discuss protons, neutrons and electrons because they are the only ones that affect chemical behavior -Charge on proton = +1.602 x 10 -19 C. Assigned a +1 charge. -Charge on an electron = -1.602 x 10 -19 C. Assigned a -1 charge. -Neutrons – Electrically neutral. Similar mass to protons. Discovered by Chadwick.

12 2.3 – Modern Atomic Theory amu = atomic mass unit. 1 amu = 1.66054 x10 -24 g. Based on 12 amu = mass of one atom of Carbon- 12. 1 amu is approximately equal to the mass of a proton/neutron. Angstrom (Å) = Unit used to measure atom size.Å 1 Angstrom = 1 x 10 -10 m

13 2.3 – Modern Atomic Theory Atomic Numbers, Mass Numbers and Isotopes Symbols: 2. Carbon - 12 1. Mass Number

14 2.3 – Atomic Number, Etc. All atoms of the same element have the same number of protons: The atomic number (Z) Atomic Number:

15 2.3 – Atomic Number, Etc. The total number of protons and neutrons in the atom. Mass Number:

16 2.3 – Atomic Number, Etc. Isotopes = Atoms of the same element with different masses. -Isotopes have different numbers of neutrons. -Isotopes have different mass numbers. 11 6 C 12 6 C 13 6 C 14 6 C Carbon-11Carbon-13 Carbon-12Carbon-14

17 2.4 – Atomic Weights Average Atomic Masses – Based on abundance of isotopes and mass of each isotope. =Atomic Weight Example 2.4 – Page 47

18 2.5 - Periodic Table The rows on the periodic table are periods. Columns are groups. Elements in the same group have similar chemical properties. Periodic Table – Arranged by atomic number with elements with similar properties in same column. Mendeleev – Organized first accepted Periodic Table

19 Periodic Table Metals are on the left side of the chart. -Low electronegativity -Good conductors of heat and electricity -Malleable, Ductile, Luster -Readily lose electrons

20 Periodic Table Nonmetals are on the right side of the periodic table (with the exception of H). -High electronegativity -Readily gain electrons

21 Periodic Table Metalloids border the stair-step line (with the exception of Al and Po). -Properties are intermediate between metals and nonmetals.

22 Groups These five groups are known by their names. 1 2 16 17 18

23 2.6 – Molecules and Molecular Compounds Chemical Formula = Shows the atoms present in a substance Molecule – Two or more atoms in a definite arrangement held by covalent bonds. Can be the same or different atoms. ex – H 2, H 2 O, C 6 H 12 O 6 Diatomic Molecule = Molecule made of only 2 atoms total Molecular Compound – Must have DIFFERENT elements ex- H 2 O is a compound, but H 2 is NOT

24 2.6 – Molecules REVIEW: Bonds – -Nonpolar Covalent – Equal sharing of electrons -No partial charges -Electronegativity of two atoms is close -Difference of electronegativities is <0.3 -Polar Covalent – Unequal sharing of electrons -Partial charges exist on atoms -Electronegativity difference is 0.3-1.7

25 2.6 - Molecules Molecular Formula = Gives actual number and type of atoms in a molecule Empirical Formula = Simplest whole number ration of atoms Example = Glucose

26 2.6 - Molecules Structural Formula – Similar to a Lewis Structure. Shows bonds, but not shape.

27 2.6 - Molecules

28 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. – Anions are negative and are formed by elements on the right side of the periodic chart. 2.7 – Ions and Ionic Compounds

29 2.7 - Ions Examples:

30 2.7 - Ions Ionic compounds (such as NaCl) are formed by ionic bonds. Contain both + and – ions. Ionic bonds are generally formed between metals (cations) and nonmetals (anions). Monatomic – one atom only with a charge Polyatomic – more than one atom with an overall charge

31 Metallic Bond Metal ions with an electron sea.

32 2.8 – Naming Inorganic Compounds Ionic Names and Formulas of Ionic Compounds: 1.Cations a.One charge only = name only Zn +2 = Zinc ion Ca +2 = Calcium ion b.Multiple charges possible = use Roman Numeral Cu +1 = Copper (I) Cu +2 = Copper (II) -Old system uses –ous / -ic c.Polyatomics – hydronium, ammonium

33 Common Cations

34 2.8 – Inorganic Compounds Ionic Names and Formulas of Ionic Compounds: 2. Anions: a.Monotomic – Change ending to –ide b.Polyatomic – MEMORIZE 8-ates and rules! c.H+ added ions = use bi- or hydro-prefix Example – bicarbonate = hydrogen carbonate CarbonateCO 3 -2 NitrateNO 3 -1 SulfateSO 4 -2 ChlorateClO 3 -1 ChromateCrO 4 -2 BromateBrO 3 -1 PhosphatePO 4 -3 IodateIO 3 -1

35 Common Anions

36 2.8 – Inorganic Compounds Ionic Ionic Compounds – Cation + Anion -Binary – Compound of 2 “atoms” -ex – NaCl, Al 2 O 3 -Ternary – Compound of 3 “atoms” -ex – Mg(NO 3 ) 2

37 2.8 – Inorganic Acids

38 2.8 – Inorganic Compounds Binary Molecular Molecular Compounds: Discrete molecular units. Contain covalent bonds. Use prefixes.

39 2.9 - Organic


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