1 chemistry second edition julia burdge lecture powerpoints jason a. kautz university of...

1

ChemistrySecond Edition

Julia Burdge

Lecture PowerPointsJason A. Kautz

University of Nebraska-Lincoln

Copyright (c) The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2Atoms, Molecules, and

Ions

Atoms, Molecules, and Ions2

2.1 The Atomic Theory2.2 The Structure of the Atom

Discovery of the Electron

RadioactivityThe Proton and the

NucleusNuclear Model of the

AtomThe Neutron

2.3 Atomic Number, Mass Number, and Isotopes2.4 The Periodic Table2.5 The Atomic Mass Scale and

Average Atomic Mass2.6 Molecules and Molecular

CompoundsMoleculesMolecular FormulasNaming Molecular

CompoundsEmpirical Formulas

2.7 Ions and Ionic CompoundsAtomic IonsPolyatomic IonsFormulas of Ionic

CompoundsNaming of Ionic

CompoundsHydratesFamiliar Inorganic

Compounds

The Atomic Theory

In 1808, John Dalton formulated a precise definition of matter that we call atoms:

1) Elements are composed of extremely small particles called atoms. All atoms of a given element are identical. The atoms of one element are different from the atoms of all other elements.

2) Compounds are composed of atoms of more than one element. In any given compound the same types of atoms are always present in the same relative numbers.

3) A chemical reaction rearranges atoms in chemical compounds; it does not create or destroy them.

2.1

The Atomic TheoryIn the reaction below:

All the oxygen molecules appear identical to one another (hypothesis 1).

The compound CO2 forms when each carbon atoms combines with two oxygen atoms (hypothesis 2).

All atoms present before the reaction are also present after the reaction (hypothesis 3).

The Atomic Theory

According to Proust’s law of definite proportions, different samples of a given compound always contain the same elements in the same mass ratio.

Sample Mass of O (g) Mass of C (g) Ratio (g O : g C)

123 g carbon dioxide 89.4 33.6 2.66:1

50.5 g carbon dioxide 36.7 13.8 2.66:1

88.6 g carbon dioxide 64.4 24.2 2.66:1

The Atomic Theory

According to the law of multiple proportions, if two elements can combine to form more than one compound with each other, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers.

12

=

12

=O to C ratio in COO to C ratio in CO2

The Structure of the Atom

Based on Dalton’s atomic theory, an atom is the basic unit of an element that can enter into a chemical combination.

2.2

The Structure of the AtomWhen metal plates are connected to a high-voltage source, the negatively charged plate, or cathode, emits an invisible ray.

The cathode ray is drawn to the anode.

Cathode(-)

Anode(+)


Cathode rays may be deflected by magnetic or electric fields.

Particles are negatively charged and are known as electrons.


J. J. Thomson determined the charge to mass ratio of an electron:

charge : mass = 1.76 × 108 C/g

C stands for coulomb which is the derived SI unit for charge


R. A. Millikan measured the charge of an electron with great precision.

1928

8

charge 1.6022 10 Cmass of an electron = = = 9.10 10 gcharge / mass 1.76 10 C/g

The Structure of the AtomThree types of decay particles produced by radioactive decay include:

Alpha (α) rays; positively charged

Beta (β) rays; negatively charge

Gamma (γ) rays; no charge; very high energy waves


Ernest Rutherford used α particles to prove the structure of atoms.

The majority of particles penetrated the gold foil undeflected.

Sometimes, α particles were deflected at a large angle.

Sometimes, α particles bounced back in the direction from which they had come.

The Structure of the AtomRutherford proposed a new model for the atom:

Positive charge is concentrated in the nucleus

The nucleus accounts for most of an atom’s mass and is an extremely dense central core within the atom

A typical atomic radius is about 100 pm

A typical nucleus has a radius of about 5 x 10–3 pm

1 pm = 1 x 10–12 m


Protons are positively charged particles found in the nucleus.

Neutrons are electronically neutral particles found in the nucleus.

Neutrons are slightly larger than protons.

Atomic Number, Mass Number, and Isotopes

Elemental symbolXZA

All atoms can be identified by the number of protons and neutrons they contain.

The atomic number (Z) is the number of protons in the nucleus

The mass number (A) is the total number of protons and neutrons.

Mass number(number of protons + neutrons)

Atomic number(number of protons)

2.3


Most elements have two or more isotopes, atoms that have the same atomic number but different mass numbers.

hydrogen

1 proton0 neutrons

deuterium

1 proton1 neutron

tritium

1 proton2 neutrons

11H

31H

21H

Isotopes of the same element exhibit similar chemical properties, forming the same types of compounds and displaying similar reactivities.


How many protons neutrons and electrons are in an atom of boron-10?

Solution:

Atomic number (Z) = 5

Mass number (A) = 10

Protons = 5 (atomic number)Electrons = 5 (same number of protons and electrons in an atom)Neutrons = 5 (mass number – atomic number; 10 – 5 = 5)

105B

The Periodic Table

The periodic table is a chart in which elements having similar chemical and physical properties are grouped together.

2.4

The Periodic Table

Elements are grouped in horizontal rows called periods.Vertical columns are called groups or families.

Elements can be categorized as metals, nonmentals or metalloids.

Metals are good conductorsof heat and electricity.

Nonmetals are poor conductors of heat orelectricity.

Metalloids have Intermediate properties.

The Periodic Table

The Atomic Mass Scale and Average Atomic Mass

The most direct way to measure molecular mass is in a mass spectrometer.

2.5

Atomic mass is the mass of an atom in atomic mass units (amu).

1 amu = 1/12 the mass of a carbon-12 atom

The atomic weight on the periodic table represents the average mass of the naturally occurring mixture of isotopes.

Average mass (C) = (0.9893)(12.00000 amu) + (0.0107)(13.003355 amu) Average mass (C) = 12.01 amu

Isotope Isotopic mass (amu)

Natural abundance (%)

12C 12.00000 98.93

13C 13.003355 1.07


The atomic mass and natural abundances of the two stable isotopes of copper are given below. Calculate the average atomic mass.

Solution:

Average mass (Cu) = (0.6917)(62.929599 amu) + (0.3083)(64.927793 amu) Average mass (Cu) = 63.55 amu

Isotope Isotopic mass (amu)

Natural abundance (%)

63Cu 62.929599 69.17

65Cu 64.927793 30.83


Molecules and Molecular Compounds

A molecule is a combination of at least two atoms in a specific arrangement held together by chemical forces (chemical bonds).

A molecule may be an element or a compound.

Diatomic molecules contain two atoms and may be either heteronuclear or homonuclear.

Polyatomic molecules contain more than two atoms.

2.6

A chemical formula denotes the composition of the substance.

A molecular formula shows the exact number of atoms of each element in a molecule.

A structural formula shows not only the elemental composition, but also the general arrangements.



Binary molecular compounds are substances that consist of just two different elements.

Nomenclature:

1) Name the first element that appears in the formula

2) Name the second element that appears in the formula, changing its ending to –ide.

Examples:

HCl hydrogen chloride

HI hydrogen iodide.

Molecules and Molecular CompoundsGreek prefixes are used to denote the number of atoms of each element present.

Examples:

SO2 sulfur dioxide

CO carbon monoxide

The prefix mono- is generally omitted for the first element.

N2O5 dinitrogen pentoxide


Name the following binary molecular compounds: (a) Cl2O

(b) SiCl4

Solution:

(b) Cl2O dichlorine monoxide

(b) SiCl4 silicone tetrachloride


Write the formula for the following compounds

(a) carbon disulfide

(b) dinitrogen trioxide

Solution:

(b) CS2

(b) N2O3


One definition of an acid is a substance that produces hydrogen ions (H+) when dissolved in water.

HCl is an example of a binary compound that is an acid when dissolved in water.

To name these types of acids:

1) remove the –gen ending from hydrogen

2) change the –ide ending on the second element to –ic.

hydrogen chloride + –ic acid → hydrochloric acid


A compound must contain at least one ionizable hydrogen atom to be an acid upon dissolving.

Examples:

HF hydrofluoric acid

HCl hydrochloric acid

HBr hydrobromic acid

HI hydroiodic acid


Organic compounds contain carbon and hydrogen, sometimes in combination with other atoms.

Hydrocarbons contain only carbon and hydrogen

The simplest hydrocarbons are called alkanes.


Many organic compounds contain groups of atoms known as functional groups.

Functional groups determine many of the chemical properties of a compound.

Reactivity typically takes place at functional groups.


Molecular formulas give the exact number of each type of element in a compound.

Empirical formulas give the simplest ratio of elements in a compound.

Compound Molecular Formula Empirical Formula

water H2O H2O

hydrogen peroxide H2O2 HO

benzene C6H6 CH


Write the empirical formula for the following compound:

C8H10N4O2

Solution:

Step 1: The elements are in an 8:10:4:2 ratio, which may be simplified to a 4:5:2:1 ratio (divide by the smallest number)

Empirical formula

C4H5N2O1

Ions and Ionic Compounds

An atomic ion or monatomic ion is one that consists of just one atom with a positive or negative charge.

The loss of one or more electrons from an atom yields a cation.

Cations have a positive charge.

Na Atom Na+ Ion

11 protons 11 protons

11 electrons 10 electrons

2.7


An anion is an ion whose net charge is negative due to an increase in the number of electrons.

Cl Atom Cl– Ion

17 protons 17 protons

17 electrons 18 electrons


Atoms can lose or gain more than one electron.


A monatomic ion is named by changing the ending of the element’s name to –ide.

Cl– is chloride O2– is oxide

Some metals can form cations of more than one possible charge.

Fe2+ : ferrous ion [Fe(II)]Fe3+ : ferric ion [Fe(III)]

Mn2+ : manganese(II) ionMn3+ : manganese(III) ionMn4+ : manganese(IV) ion


Ions that consist of a combination of two or more atoms are called polyatomic ions.


Formulas for ionic compounds are generally empirical formulas.

Ionic compounds are electronically neutral.


Al3+ O2–

Al2O3

In order for ionic compounds to be electronically neutral, the sum of the charges on the cation and anion in each formula must be zero.

Aluminum oxide:

Sum of charges: 2(+3) + 3(–2) = 0


To name ionic compounds:

1) Name the cation .

omit the word ion

use a Roman numeral if the cation can have more than one charge

2) Name the anion

omit the word ion

Examples:

NaCN sodium cyanide

FeCl2 iron(II) chloride

FeCl3 iron(III) chloride


Name the following ionic compounds

(a) Na2SO4

(b) Cu(NO3)2

(c) CuNO3

Solution:

(d) sodium sulfate

(e) copper(II) nitrate

(f) copper(I) nitrate


Oxoanions are polyatomic anions that contain one or more oxygen atoms and one atom of another element.

Starting with the oxoanions whose names end in –ate, the following can be used to name other oxoanions in the same family:

1) The ion with one more O atom than the –ate ion is called per . . . ate ion. (ClO3

- is chlorate; ClO4- is perchlorate)

2) The ion with one less O atom than the –ate anion is called the

–ite ion. (ClO2- is chlorite)

3) The ion with two fewer O atoms than the –ate ion is called the

hypo . . . ite ion (ClO- is hypochlorite)


Oxoacids are acids that contain oxoanions.

1) An acid based on the –ate ion is called . . . ic acid

HClO3 is chloric acid

2) An acid based on the –ite ion is called . . . ous acid

HClO2 is chlorous acid

3) Prefixes in oxoanion names are retained in the name of the acid.

HClO4 is perchloric acidHClO is hypochlorous acid


Hydrates are compounds that have a specific number of water molecules within their solid structure.

CuSO4 • 5 H2O copper(II) sulfate pentahydrate

Chapter Summary: Key Points2

The Atomic TheoryDiscovery of the ElectronRadioactivityThe Proton and the NucleusNuclear Model of the AtomThe NeutronThe Periodic TableAverage Atomic MassMolecular FormulasNaming Molecular Compounds

Empirical FormulasAtomic IonsPolyatomic IonsFormulas of Ionic CompoundsNaming Ionic CompoundsHydratesInorganic Compounds

1 chemistry second edition julia burdge lecture powerpoints jason a. kautz university of...

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