Chapter 2

Atoms, Molecules, and Ions

AP*

AP Learning Objectives

! LO 1.1 The student can justify the observation that the ratio of the masses of the constituent elements in any pure sample of that compound is always identical on the basis of the atomic molecular theory. (Sec 2.2)

! LO 1.17 The student is able to express the law of conservation of mass quantitatively and qualitatively using symbolic representations and particulate drawings. (Sec 2.2, 2.3)

! LO 2.17 The student can predict the type of bonding present between two atoms in a binary compound based on position in the periodic table and the electronegativity of the elements. (Sec 2.6, 2.7)

! LO 3.5 The student is able to design a plan in order to collect data on the synthesis or decomposition of a compound to confirm the conservation of matter and the law of definite proportions. (Sec 2.2)

! LO 3.6 The student is able to use data from synthesis or decomposition of a compound to confirm the conservation of

AP Learning Objectives

chapter 2

! IONS, ATOMS AND MOLECULES

AP Learning Objectives

! Alchemy ! obsessed with the idea of turning cheap

metals into gold.

AP Learning Objectives

Robert Boyle

! Robert Boyle was the first “chemist”. ! Performed quantitative experiments.

! The skeptical chemist 1661

AP Learning Objectives

Boyle’s investigations

! pressure and volume of air !

AP Learning Objectives

Fundamental Chemical Laws! Anthonie Lavosier

(1743-1794). ! explained combustion ! invested in a tax

collecting firm and married the daughter of one of the executives

AP Learning Objectives

AP Learning Objectives

French Revolution 1794

! Executed in the guillotine as an enemy of the people

AP Learning Objectives

AP Learning Objectives

John Dalton 1766-1844

! Dalton’s law ! handicaps: color blind ! atom model

AP Learning Objectives

Joseph Proust 1754-1826! a given compound

always contains the same proportion of elements by mass. !

! law of definite proportions

"Atoms combine in whole number ratios, so their proportion by mass will always be the same.

"Example: H2O is always made up of 2 atoms of H and one atom of O.

The ratio of O to H in water is always 16:2 or 8:1.

Example:

"KCl always contains one atom of K for every one atom of Cl

" In KCl, potassium and chlorine always have a ratio of “39.09 to 35.45” or “1.1 to 1” by mass.

Practice Problem 1

In the carbon compounds ethane (C2H6) and ethene (C2H4), what is the lowest whole number ratio of H atoms that react with the same number of C atoms?

!!Answer: 3:2

AP Learning Objectives

Law of definite proportions

AP Learning Objectives

Dalton’s experiments

! Analyzed the relative masses of elements in different compounds

AP Learning Objectives

Mass of Nitrogen that combines with 1g of oxygen

compound A 1.750g

compound B .8750g

compound C .4375g

AP Learning Objectives

A/B 1.750/.8750 2/1

B/C .8750/.4375 2/1

A/C 1.750/.4375 4/1

ILLUSTRATING THE LAW OF MULTIPLE PROPORTIONS

AP Learning Objectives

DALTON’S ATOMIC THEORY

AP Learning Objectives

Dalton prepared the first table of atomic masses... he called them atomic weights.

AP Learning Objectives

Joseph Gay Lussac, French 1778-1850

! same conditions of temperature and pressure.

! experimented with the volumes of gases that reacted with each other.

AP Learning Objectives

Gay Lussac

AP Learning Objectives

Gay Lussac

! 2 volumes of hydrogen react with 1 volume of oxygen to form 2 volumes of gaseous water.

Section 2.3 Dalton’s Atomic Theory

Representing Gay—Lussac’s Results

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Section 2.3 Dalton’s Atomic Theory

Representing Gay—Lussac’s Results

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Section 2.3 Dalton’s Atomic Theory

Representing Gay—Lussac’s Results

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Section 2.3 Dalton’s Atomic Theory

Representing Gay—Lussac’s Results

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Section 2.3 Dalton’s Atomic Theory

Lorenzo Romano Amedeo Carlo Avogadro di Quaregna e di Cerreto

AP Learning Objectives

Avogadro

! At the same temperature and pressure, equal volumes of different gases contain the same number of particles.

AP Learning Objectives

Amadeo Avogadro

! Avogadro’s number !

!

!! described the number of particles in a

substance.

AP Learning Objectives

J. J. Thomson

! described the electron ! the charge of an electron ! behavior electron

AP Learning Objectives

! Discovered the charge to mass ratio of an electron.

! e/m= ! e represents the charge of the electron ! m represents the electron mass in grams

-1.76 x 10^(8) C/g

TextTextText

Section 2.4 Early Experiments to Characterize the AtomRobert Millikan (1909)

! Performed experiments involving charged oil drops.

! Determined the magnitude of the charge on a single electron.

! Calculated the mass of the electron ! (9.11 × 10-31 kg).

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Section 2.4 Early Experiments to Characterize the AtomHenri Becquerel (1896)

! Discovered radioactivity by observing the spontaneous emission of radiation by uranium.

! Three types of radioactive emission exist: ! Gamma rays (ϒ) – high energy light ! Beta particles (β) – a high speed electron ! Alpha particles (α) – a particle with a 2+

charge

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Section 2.4 Early Experiments to Characterize the AtomRutherford

! he passed an alpha particle (2 protons) through a sheet of metal foil. !

!

!

!

!! and thus, the atom is mostly empty

AP Learning Objectives

Rutherford

! mostly empty spaced

AP Learning Objectives

Bohr Model

! energy levels

AP Learning Objectives

Cloud Model

! We can never determine where the electron is located.

! know where it spends most of the time.

AP Learning Objectives

what we know

! proton + : size and charge +1 ! neutron: size and charge (0) !

! electron: size and charge -1

Section 2.1 The Early History of Chemistry

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! Greeks were the first to attempt to explain why chemical changes occur.

! Alchemy dominated for 2000 years. ! Several elements discovered. ! Mineral acids prepared.

! Robert Boyle was the first “chemist”. ! Performed quantitative experiments. ! Developed first experimental definition

of an element.

Early History of Chemistry

Section 2.2 Fundamental Chemical Laws

AP Learning Objectives, Margin Notes and ! Learning Objectives ! LO 1.1 The student can justify the observation that the ratio of the masses of the constituent

elements in any pure sample of that compound is always identical on the basis of the atomic molecular theory.

! LO 1.17 The student is able to express the law of conservation of mass quantitatively and qualitatively using symbolic representations and particulate drawings.

! Additional AP References ! LO 3.5 (see APEC Lab 7, " Hydrates and Thermal Decomposition") ! LO 3.6 (see APEC Lab 7, " Hydrates and Thermal Decomposition")

Section 2.2 Fundamental Chemical Laws

Three Important Laws

! Law of conservation of mass (Lavoisier): ! Mass is neither created nor destroyed in a

chemical reaction. ! Law of definite proportion (Proust):

! A given compound always contains exactly the same proportion of elements by mass.

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Section 2.2 Fundamental Chemical Laws

Three Important Laws (continued)

! Law of multiple proportions (Dalton): ! When two elements form a series of

compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers.

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Section 2.3 Dalton’s Atomic TheoryAP Learning Objectives, Margin Notes and References! Learning Objectives ! LO 1.17 The student is able to express the law of conservation of mass quantitatively and

qualitatively using symbolic representations and particulate drawings.

Section 2.3 Dalton’s Atomic Theory

Dalton’s Atomic Theory (1808)

! Each element is made up of tiny particles called atoms.

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Section 2.3 Dalton’s Atomic Theory

Dalton’s Atomic Theory (continued)

! The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways.

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Section 2.3 Dalton’s Atomic Theory

Dalton’s Atomic Theory (continued)

! Chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative numbers and types of atoms.

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Section 2.3 Dalton’s Atomic Theory

Dalton’s Atomic Theory (continued)

! Chemical reactions involve reorganization of the atoms—changes in the way they are bound together.

! The atoms themselves are not changed in a chemical reaction.

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Section 2.3 Dalton’s Atomic Theory

Gay-Lussac and Avogadro (1809—1811)

! Gay—Lussac ! Measured (under same conditions of T and

P) the volumes of gases that reacted with each other.

! Avogadro’s Hypothesis ! At the same T and P, equal volumes of

different gases contain the same number of particles. ! Volume of a gas is determined by the

number, not the size, of molecules.12

Section 2.3 Dalton’s Atomic Theory

Representing Gay—Lussac’s Results

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Section 2.3 Dalton’s Atomic Theory

Representing Gay—Lussac’s Results

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Section 2.3 Dalton’s Atomic Theory

Representing Gay—Lussac’s Results

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Section 2.3 Dalton’s Atomic Theory

Representing Gay—Lussac’s Results

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Section 2.4 Early Experiments to Characterize the AtomJ. J. Thomson (1898—1903)

! Postulated the existence of negatively charged particles, that we now call electrons, using cathode-ray tubes.

! Determined the charge-to-mass ratio of an electron.

! The atom must also contain positive particles that balance exactly the negative charge carried by electrons.

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Section 2.4 Early Experiments to Characterize the AtomCathode-Ray Tube

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Section 2.4 Early Experiments to Characterize the AtomRobert Millikan (1909)

! Performed experiments involving charged oil drops.

! Determined the magnitude of the charge on a single electron.

! Calculated the mass of the electron ! (9.11 × 10-31 kg).

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Section 2.4 Early Experiments to Characterize the AtomMillikan Oil Drop Experiment

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Section 2.4 Early Experiments to Characterize the AtomHenri Becquerel (1896)

! Discovered radioactivity by observing the spontaneous emission of radiation by uranium.

! Three types of radioactive emission exist: ! Gamma rays (ϒ) – high energy light ! Beta particles (β) – a high speed electron ! Alpha particles (α) – a particle with a 2+

charge

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Section 2.4 Early Experiments to Characterize the AtomErnest Rutherford (1911)

! Explained the nuclear atom. ! The atom has a dense center of positive charge

called the nucleus. ! Electrons travel around the nucleus at a large

distance relative to the nucleus.

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Section 2.5 The Modern View of Atomic Structure: An Introduction

! The atom contains: ! Electrons – found outside the nucleus;

negatively charged. ! Protons – found in the nucleus; positive

charge equal in magnitude to the electron’s negative charge.

! Neutrons – found in the nucleus; no charge; virtually same mass as a proton.

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Section 2.5 The Modern View of Atomic Structure: An Introduction

! The nucleus is: ! Small compared with the overall size of the

atom. ! Extremely dense; accounts for almost all of

the atom’s mass.

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Section 2.5 The Modern View of Atomic Structure: An Introduction

Nuclear Atom Viewed in Cross Section

Section 2.5 The Modern View of Atomic Structure: An IntroductionIsotopes

! Atoms with the same number of protons but different numbers of neutrons.

! Show almost identical chemical properties; chemistry of atom is due to its electrons.

! In nature most elements contain mixtures of isotopes.

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Section 2.5 The Modern View of Atomic Structure: An Introduction

Two Isotopes of Sodium

Section 2.5 The Modern View of Atomic Structure: An Introduction

! Isotopes are identified by: ! Atomic Number (Z) – number of protons ! Mass Number (A) – number of protons plus

number of neutrons

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Section 2.5 The Modern View of Atomic Structure: An Introduction

A certain isotope X contains 23 protons and 28 neutrons.

! What is the mass number of this isotope? ! Identify the element. !

Mass Number = 51 Vanadium

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EXERCISE!

Section 2.6 Molecules and IonsAP Learning Objectives, Margin Notes and References! Learning Objectives ! LO 2.17 The student can predict the type of bonding present between two atoms in a binary

compound based on position in the periodic table and the electronegativity of the elements.

Section 2.6 Molecules and Ions

Chemical Bonds

! Covalent Bonds ! Bonds form between atoms by sharing

electrons. ! Resulting collection of atoms is called a

molecule.

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Section 2.6 Molecules and Ions

Covalent Bonding

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Section 2.6 Molecules and Ions

types of bonding animation link

! http://www.kentchemistry.com/links/bonding/bondingflashes/bond_types.swf

Section 2.6 Molecules and Ions

Chemical Bonds

! Ionic Bonds ! Bonds form due to force of attraction

between oppositely charged ions. ! Ion – atom or group of atoms that has a net

positive or negative charge. ! Cation – positive ion; lost electron(s). ! Anion – negative ion; gained electron(s).

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Section 2.6 Molecules and Ions

Molecular vs Ionic Compounds

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Section 2.6 Molecules and Ions

A certain isotope X+ contains 54 electrons and 78 neutrons.

!! What is the mass number of this isotope? !

133

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EXERCISE!

Section 2.6 Molecules and Ions

Which of the following statements regarding Dalton’s atomic theory are still believed to be true? !I. Elements are made of tiny particles called atoms. II. All atoms of a given element are identical. III. A given compound always has the same relative numbers and types of atoms. IV. Atoms are indestructible.

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CONCEPT CHECK!

Section 2.7 An Introduction to the Periodic TableAP Learning Objectives, Margin Notes and References! Learning Objectives ! LO 2.17 The student can predict the type of bonding present between two atoms in a binary

compound based on position in the periodic table and the electronegativity of the elements.

Section 2.7 An Introduction to the Periodic TableThe Periodic Table http://www.dnatube.com/video/11465/! Metals vs. Nonmetals ! Groups or Families – elements in the same

vertical columns; have similar chemical properties

! Periods – horizontal rows of elements

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Section 2.7 An Introduction to the Periodic TableThe Periodic Table

Section 2.7 An Introduction to the Periodic Table

Dimitri Mendelev

• He was perpetually in trouble with the Russian government and the Russian Orthodox Church, but he was brilliant never-the-less.

Section 2.7 An Introduction to the Periodic Table

Section 2.7 An Introduction to the Periodic Table

trends in the periodic table

! http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/atomic4.swf

Section 2.7 An Introduction to the Periodic Table

Groups or Families

! Table of common charges formed when creating ionic compounds.

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Group or Family ChargeAlkali Metals (1A) 1+

Alkaline Earth Metals (2A) 2+Halogens (7A) 1–

Noble Gases (8A) 0

Section 2.8 Naming Simple Compounds

Naming Compounds

! Binary Compounds ! Composed of two elements ! Ionic and covalent compounds included

! Binary Ionic Compounds ! Metal—nonmetal

! Binary Covalent Compounds ! Nonmetal—nonmetal

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Section 2.8 Naming Simple CompoundsBinary Ionic Compounds (Type I) cation and anion1. The cation is always named first and the anion

second. 2. A monatomic cation takes its name from the

name of the parent element. 3. A monatomic anion is named by taking the

root of the element name and adding –ide.

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Section 2.8 Naming Simple Compounds

Binary Ionic Compounds (Type I)

! Examples: KCl Potassium chloride ! MgBr2 Magnesium bromide ! CaO Calcium oxide

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Section 2.8 Naming Simple Compounds

Binary Ionic Compounds (Type II)

! Metals in these compounds form more than one type of positive ion.

! Charge on the metal ion must be specified. ! Roman numeral indicates the charge of the

metal cation. ! Transition metal cations usually require a

Roman numeral. ! Elements that form only one cation do not

need to be identified by a roman numeral.44

Section 2.8 Naming Simple Compounds

Binary Ionic Compounds (Type II)

! Examples: CuBr Copper(I) bromide ! FeS Iron(II) sulfide ! PbO2 Lead(IV) oxide

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Section 2.8 Naming Simple Compounds

Polyatomic Ions

! Must be memorized (see Table 2.5 on pg. 65 in text).

! Examples of compounds containing polyatomic ions:

NaOH Sodium hydroxide Mg(NO3)2 Magnesium nitrate (NH4)2SO4 Ammonium sulfate

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Section 2.8 Naming Simple Compounds

Formation of Ionic Compounds

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Section 2.8 Naming Simple Compounds

Binary Covalent Compounds (Type III)

! Formed between two nonmetals. 1. The first element in the formula is named first,

using the full element name. 2. The second element is named as if it were an

anion. 3. Prefixes are used to denote the numbers of

atoms present. 4. The prefix mono- is never used for naming the

first element.48

Section 2.8 Naming Simple Compounds

Prefixes Used to Indicate Number in Chemical Names

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Section 2.8 Naming Simple Compounds

Binary Covalent Compounds (Type III)

! Examples: CO2 Carbon dioxide ! SF6 Sulfur hexafluoride ! N2O4 Dinitrogen tetroxide

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Section 2.8 Naming Simple Compounds

Flowchart for Naming Binary Compounds

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Section 2.8 Naming Simple CompoundsOverall Strategy for Naming Chemical Compounds

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Section 2.8 Naming Simple Compounds

Acids

! Acids can be recognized by the hydrogen that appears first in the formula—HCl.

! Molecule with one or more H+ ions attached to an anion.

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Section 2.8 Naming Simple Compounds

Acids

! If the anion does not contain oxygen, the acid is named with the prefix hydro– and the suffix –ic.

! Examples: HCl Hydrochloric acid HCN Hydrocyanic acid H2S Hydrosulfuric acid

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Section 2.8 Naming Simple Compounds

Acids

! If the anion does contain oxygen: ! The suffix –ic is added to the root name if

the anion name ends in –ate. ! Examples: HNO3 Nitric acid H2SO4 Sulfuric acid HC2H3O2 Acetic acid

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Section 2.8 Naming Simple Compounds

Acids

! If the anion does contain oxygen: ! The suffix –ous is added to the root name if

the anion name ends in –ite. ! Examples: HNO2 Nitrous acid H2SO3 Sulfurous acid HClO2 Chlorous acid

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Section 2.8 Naming Simple Compounds

Flowchart for Naming Acids

Section 2.8 Naming Simple Compounds

Which of the following compounds is named incorrectly? !a) KNO3 potassium nitrate b) TiO2 titanium(II) oxide c) Sn(OH)4 tin(IV) hydroxide d) PBr5 phosphorus pentabromide e) CaCrO4 calcium chromate

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EXERCISE!