Atoms, Molecules and Ions. Chapter 2
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1 Atoms, Molecules and Ions Chapter 2
2 2.1 The Atomic Theory of Matter Democritus [ BCE] Described tiny, indivisible particles Called them atomos Differed from Aristotle 17th century - idea of atoms returned to Europe
3 2.1 The Atomic Theory of Matter John Dalton [ ] Developed beginnings of modern atomic theory Wrote postulates to explain several laws Law of definite proportions [postulate 4] Law of conservation of mass [postulate 3] Developed law of multiple proportions
4 Dalton s Atomic Theory 1. Each element is composed of extremely small particles called atoms. 2. All atoms of a given element are identical, but the atoms of one element are different from the atoms of all other elements. 3. Atoms of one element cannot be changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions. 4. Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.
5 Give It Some Thought Compound A contains g of oxygen per gram of carbon, whereas compound B contains g of oxygen per gram of carbon. a. What chemical law do these data illustrate? b. If compound A has an equal number of carbon and oxygen atoms, what can we conclude about the composition of compound B?
6 2.2 The Discovery of Atomic Structure J.J. Thomson [ ] Studied cathode rays, which were known to be negatively charged Noticed it didn t matter what the material was for the cathode Credited with the discovery of the electron [paper published in 1897] Measured mass to charge ratio of electrons
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8 2.2 The Discovery of Atomic Structure
9 2.2 The Discovery of Atomic Structure Robert Millikan [ ] Measured the charge of an electron
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11 2.2 The Discovery of Atomic Structure Ernest Rutherford [ ] Shot α particles at thin gold foil Most passed through Some deflected at sharp angles Developed nuclear model nucleus mostly empty space
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13 2.2 The Discovery of Atomic Structure Protons Found in 1919 by Rutherford Neutrons Found in 1932 by James Chadwick
14 HOMEWORK 2.11, 2.12, 2.13, 2.17
15 2.3 The Modern View of Atomic Structure Physicists found more particles [quarks] that make up the protons and neutrons Chemists only focus on protons, neutrons, and electrons for chemical behavior
16 2.3 The Modern View of Atomic Structure Particle Charge Mass (amu) Location Proton Positive (1+) Nucleus Neutron Electron None (neutral) Negative (1- ) Nucleus x 10-4 Electron Cloud
17 2.3 The Modern View of Atomic Structure Units developed to talk about atoms atomic mass unit (amu) 1 amu = x g angstrom (Å) 1Å = 1 x m
18 2.3 The Modern View of Atomic Structure Atomic Number Number of protons in nucleus Mass Number Number of protons + neutrons in nucleus NOT FOUND ON PT Notations
19 2.3 The Modern View of Atomic Structure Isotopes Atoms of same element with different mass numbers Same protons, different neutrons
20 Atomic Number Practice 1. How many protons, neutrons, and electrons are in an atom of 197 Au? 2. How many protons, neutrons, and electrons are in an atom of strontium-90?
21 Atomic Number Practice 3. Magnesium has three isotopes with mass numbers 24, 25, and 26. Write the symbol notation for all three isotopes. How many neutrons does each isotope contain?
22 2.4 Atomic Weights average atomic mass called atomic weight decimal number on periodic table weighted average of the masses of the naturally occurring isotopes for the element.
23 2.4 Atomic Weights Example with carbon: 98.93% carbon-12 (12 amu) 1.07% carbon-13 ( amu)
24 Atomic Weight Practice Indium has two naturally occurring isotopes. One has a mass of amu and an abundance of 4.29%. The other has a mass of amu. Calculate the atomic weight of indium.
25 HOMEWORK 2.2, 2.4, 2.19, 2.21, 2.22, 2.27, 2.28, 2.29, 2.31, 2.35
26 2.5 The Periodic Table Developed in 1869 Arrangement of the elements in order of increasing atomic number, with elements having similar properties placed in vertical columns
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28 2.5 The Periodic Table Group Names: group 1 (1A) - alkali metals group 2 (2A) - alkaline-earth metals groups 3-12 (1B-8B) group 16 (6A) - chalcogens group 17 (7A) - halogens group 18 (8A) - noble gases
29 2.5 The Periodic Table Metals left of staircase luster, conductors, solids at room temp Nonmetals right of staircase some gas, some solid, one liquid Metalloids lie on staircase
30 Periodic Table Practice 1. Which two of these elements would you expect to be most similar to each other: B, Ca, F, He, Mg, P? 2. Locate sodium and bromine in the periodic table. Give the atomic number of each and classify each as a metal, metalloid, or nonmetal.
31 2.6 Molecules and Molecular Compounds Atom is the smallest part of an element Only noble gases exist as atoms in nature Most matter is composed of molecules or ions
32 2.6 Molecules and Molecular Compounds Diatomic molecules molecule made up of two of the same element H 2, O 2, F 2, Br 2, I 2, N 2, Cl 2 Molecular compounds Exist as molecules [covalent bonds] Contain more than one type of atom
33 2.6 Molecules and Molecular Compounds Types of formula 1. Molecular formula 2. Empirical formula 3. Structural formula
34 2.6 Molecules and Molecular Compounds Picturing molecules 1. ball-and-stick models 2. space filling models
35 2.6 Molecules and Molecular Compounds
36 HOMEWORK 2.3, 2.41, 2.43, 2.49, 2.52
37 2.7 Ions and Ionic Compounds Ion Atom that has lost or gained electrons Cation Anion Positive ion Lost electron(s) Negative ion Gained electron(s)
38 Ion Practice Writing symbols for ions Give the chemical symbol, including superscript mass number, for the ion with 22 protons, 26 neutrons, and 19 electrons. Give the chemical symbol, including superscript mass number, for the ion of sulfur that has 16 neutrons and 18 electrons.
39 2.7 Ions and Ionic Compounds Polyatomic Ions groups of atoms covalently bonded, but carry a net positive or negative charge
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42 2.7 Ions and Ionic Compounds Ionic Compounds Transfer of electrons Composed of cations and anions Combination of metal and nonmetal Create a three-dimensional arrangement of ions
43 2.7 Ions and Ionic Compounds Formulas for ionic compounds always empirical formula [lowest ratio] CROSS CHARGES Examples:
44 2.8 Naming Inorganic Compounds Ionic compounds Cations Name of the metal If a transition metal, must include Roman numeral Name of the polyatomic
45 2.8 Naming Inorganic Compounds Ionic compounds Anions Change ending of element name to -ide Name polyatomic ion
46 Ionic Compound Practice 1. Name: a. CaF 2 b. FeCl 3 c. CoNO 3 d. Na 2 SO 4 2. Write the formulas: a. Potassium oxide b. Copper(II) nitride c. Aluminum carbonate d. Chromium(II) chlorate
47 2.8 Naming Inorganic Compounds Binary Covalent Compounds Name first element, with correct prefix No prefix if only one Name second element, with correct prefix, changing ending to -ide
48 Covalent Compound Practice 1. Name the following: a. N 2 O b. PCl 5 c. SO 3 d. P 4 O Write the formula. a. Silicon tetrachloride b. Ammonia c. Arsenic trifluoride
49 2.8 Naming Inorganic Compounds Naming Acids Binary hyrdo- -ic acid Oxyacid -ate polyatomics go to -ic -ite polyatomics go to -ous
50 Acid Practice 1. Name the following acids: a. HCl b. H 2 CO 3 c. HClO 2. Write formulas for the following acids: a. Hydrobromic acid b. Nitric acid c. Sulfurous acid
51 HOMEWORK 2.6, 2.7, 2.55, 2.57, 2.60, 2.62, 2.63, 2.65, 2.67, 2.71, 2.73, 2.75
52 2.9 Some Simple Organic Compounds Organic Chemistry Study of compounds of carbon
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54 2.9 Some Simple Organic Compounds Alkanes hydrocarbons Alcohols Replace a hydrogen with -OH
55 HOMEWORK 2.81, 2.83
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Basic Concepts of Chemistry Notes for Students [Chapter 2, page 1] D J Weinkauff - Nerinx Hall High School
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Basic Concepts of Chemistry Notes for Students [Chapter 2, page 1] Chapter 2 Chemical Foundations: Elements, Atoms and Ions Section 2-1: The Elements Ultimately all substances in the universe can be broken
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