AP PPT ch 2 copy
Transcript of AP PPT ch 2 copy
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.
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Robert Boyle
! Robert Boyle was the first “chemist”. ! Performed quantitative experiments.
! The skeptical chemist 1661
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Boyle’s investigations
! pressure and volume of air !
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Fundamental Chemical Laws! Anthonie Lavosier
(1743-1794). ! explained combustion ! invested in a tax
collecting firm and married the daughter of one of the executives
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French Revolution 1794
! Executed in the guillotine as an enemy of the people
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John Dalton 1766-1844
! Dalton’s law ! handicaps: color blind ! atom model
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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
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Dalton’s experiments
! Analyzed the relative masses of elements in different compounds
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Mass of Nitrogen that combines with 1g of oxygen
compound A 1.750g
compound B .8750g
compound C .4375g
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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
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DALTON’S ATOMIC THEORY
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Dalton prepared the first table of atomic masses... he called them atomic weights.
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Joseph Gay Lussac, French 1778-1850
! same conditions of temperature and pressure.
! experimented with the volumes of gases that reacted with each other.
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Gay Lussac
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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.
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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
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Bohr Model
! energy levels
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Cloud Model
! We can never determine where the electron is located.
! know where it spends most of the time.
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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!