Chemistry: The Central Science, 14e

Chemistry: The Central Science

Fourteenth Edition

Chapter 2

Atoms, Molecules, and Ions

Copyright © 2018, 2015, 2012 Pearson Education, Inc. All Rights Reserved

Atomic Theory of Matter (Dalton’s)

Basis of theory (Dalton’s Atomic Theory)

are a couple of laws.

– The law of constant composition

– The law of conservation of mass

– The law of multiple proportions

Law of Constant Composition

• (already went over in Chapter 1)

• Compounds have a definite composition. That means

that the relative number of atoms of each element in the

compound is the same in any sample. (ratio of # of

atoms of each element in a compound is the same for

the same compound)

• Example: H2O, always has 2 parts H to 1 part O

Law of Conservation of Mass: Mass is neither created

nor destroyed in chemical reactions.

Aqueous solutions of mercury(II) nitrate and potassium

iodide will react to form a precipitate of mercury(II) iodide

and aqueous potassium iodide.

HgI2(s) + 2KNO3(aq)Hg(NO3)2(aq) + 2KI(aq)

4.55 g + 2.02 g = 6.57 g

3.25 g + 3.32 g = 6.57 g

Law of Multiple Proportions

Law of Multiple Proportions: Elements can

combine in different ways to form different

substances, whose mass ratios are small whole-

number multiples of each other. (bc combining # of

atoms)

Nitrogen monoxide: 7 grams nitrogen per

8 grams oxygen (NO)

Nitrogen dioxide: 7 grams nitrogen per

16 grams oxygen (NO2)

The Law of Multiple Proportions

Law of Multiple Proportions: Elements can

combine in different ways to form different

substances, whose mass ratios are small whole-

number multiples of each other. (bc combining

discrete atoms not something like sugar and sand)

Postulates of Dalton’s Atomic Theory (1 of 4)

1) Each element is composed of extremely small

particles called atoms.


2) All atoms of an element are the same to atoms of the

same element. Atoms of different elements are

different. (C is always C, C is not N)


3) Atoms are neither created nor destroyed in chemical

reactions. Atoms are just rearranged in chemical

reactions.


4) Atoms of more than one element combine to form

compounds; a given compound always has the

same relative number and kind of atoms.

(Example: 2 H always combines with 1 O atom to

make the compound H2O)

Discovery of Subatomic Particles

Atom itself was made up of smaller particles.

– Electrons and cathode rays

– Radioactivity

– Nucleus, protons, and neutrons

The Electron (Cathode Rays)

• Streams of negatively charged particles were found to

emanate from cathode tubes, causing fluorescence.

• J. J. Thomson is credited with their discovery (1897).

The Electron

Thomson measured the charge/mass ratio of the

electron to be 81.76 10 coulombs/gram (C/g).

Millikan Oil-Drop Experiment (Electrons)

• Once the charge/mass ratio of the electron was

known, determination of either the charge or the mass

of an electron would yield the other.

• Robert Millikan determined the charge on the electron

in 1909.

Radioactivity (2 of 2)

• Three types of radiation were discovered by Ernest

Rutherford:

– particles (positively charged)

– particles (negatively charged, like electrons)

– rays (uncharged)

Discovery of the Nucleus

Ernest Rutherford shot particles at a thin sheet of

foil and observed the pattern of scatter of the particles.

Subatomic Particles

• Protons (+1) and electrons (−1) have a charge; neutrons

are neutral.

• Protons and neutrons have essentially the same mass

(relative mass 1). The mass of an electron is so small we

ignore it (relative mass 0).

• Protons and neutrons are found in the nucleus; electrons

travel around the nucleus.

Table 2.1 Comparison of the Proton, Neutron, and Electron

Particle Charge Mass (amu)

Proton Positive (1+) 1.0073

Neutron None (neutral) 1.0087

Electron Negative (1−)5.486 times 10 to the negative fourth

− 45.486 10

Atomic Numbers

Carbon-14:

C14

6

Atomic number

Mass number

Carbon-12:

C12

6

Atomic number

Mass number

6 protons

6 electrons

8 neutrons

6 protons

6 electrons

6 neutrons

A

Z

E

isotopes

Element symbol

(not periodic table)

Mass # = # Proton

+ # neutrons

Atomic Numbers

Atomic Number (Z): Number of protons in an atom’s

nucleus, equivalent to the number of electrons around an

atom’s nucleus

Mass Number (A): The sum of the number of protons

and the number of neutrons in an atom’s nucleus

Isotope: Atoms with identical atomic numbers but

different mass numbers (same # proton, different #

neutrons)

Isotopes

• Isotopes are atoms of the same element with different

masses.

• Isotopes have different numbers of neutrons, but the same

number of protons.

• The table below lists four isotopes for carbon.

Table 2.2 Some Isotopes of Carbona

Symbol Number of Protons Number of Electrons Number of Neutrons

Super 11 C 6 6 5

Super 12 C 6 6 6

Super 13 C 6 6 7

Super 14 C 6 6 8

a Almost 99% of the carbon found in nature is 12C.

11C

12C

13C

14C

HW 2.1(a) How many protons and neutrons are in

the nucleus of each of the following atoms?

(b) In a neutral atom of each element, how

many electrons are present? (look in

periodic table)

– (neutral atoms have the same # of protons as

electrons)

1. 79Br

2. 81Br

3. 239Pu

4. 133CsCopyright © Cengage Learning. All rights reserved 21

HW 2.1

• How many protons and neutrons are in the

nucleus of each of the following atoms?

– In a neutral atom of each element, how many

electrons are present ?

1. 79Br

2. 81Br

3. 239Pu

4. 133Cs

35 p, 44 n, 35 e

35 p, 46 n, 35 e

94 p, 145 n, 94 e

55 p, 78 n, 55 e

Copyright © Cengage Learning. All rights reserved 22

Periodic Table

• The periodic table is a systematic organization of the elements.

• Elements are arranged in order of atomic number.

• Accessible periodic table: media.pearsoncmg.com

https://media.pearsoncmg.com/bc/bc_0media_chem/periodictable/table.html

Atomic Masses

The mass of 1 atom of carbon-12 is defined to be 12 amu.

Atomic Mass: The weighted average of the isotopic

masses of the element’s naturally occurring isotopes

C12

6

periodic table Symbol

for

elements

HW 2.2 Calculation of Average Atomic Mass from

isotopic mass and % composition example: One

isotope of gallium has atomic mass of 68.926 amu

(atomic mass A) & makes up 60.3 % of natural gallium.

Ga also has another isotope with atomic mass of

70.925 (atomic mass B) at 39.7% What is the

average atomic mass of gallium based on the isotopic

masses ? (next slide has answer) ( 1 = 0.603 + 0.397

so if 60.3% is 68.926, then how would you do it ?)

HW 2.2 Calculation of Average Atomic Mass from

isotopic mass and % composition example: One

isotope of gallium has atomic mass of 68.926 amu

(atomic mass A) & makes up 60.3 % of natural gallium.

Ga also has another isotope with atomic mass of

70.925 (atomic mass B) at 39.7% What is the

average atomic mass of gallium based on the isotopic

masses ?

average atomic mass =

isotope A isotope B

[(fraction )(atomic mass)] + [(fraction)(atomic mass)]

(0.603) (68.926) + (0.397) (70.925) = 69.7

41.562378 + 28.157225

Organization of the Periodic Table

• The rows on the periodic table are called periods.

• Columns are called groups.

• Elements in the same group have similar chemical

properties.

Groups

Table 2.3 Names of Some Groups in the Periodic Table

Group Name Elements

1A Alkali metals Li, Na, K, Rb, Cs, Fr

2A Alkaline earth metals Be, Mg, Ca, Sr, Ba, Ra

6A Chalcogens O, S, Se, Te, Po

7A Halogens F, Cl, Br, I, At

8A Noble gases He, Ne, Ar, Kr, Xe, Rn

These five groups are known by their names.

Periodic Table (1 of 3)

• Metals are on the left side of the periodic table.

• Some properties of metals include

– Shiny luster

– Conducting heat and electricity

– Solids (except mercury)


• Nonmetals are on the right side of the periodic table

(they include H).

• They can be solid (like carbon), liquid (like bromine),

or gas (like neon) at room temperature.


• Elements on the

steplike line are

metalloids

(except Al, Po,

and At).

• Their properties

are sometimes

like metals and

sometimes like

nonmetals. End 1/21 R

Section 11.4

The Periodic Table –Divided into Periods and Groups

group

period

Chemical Formulas

• The subscript to the right of the

symbol of an element tells the number

of atoms of that element in one

molecule of the compound.

• Molecular compounds (covalent

compounds) are composed of

molecules and almost always contain

only nonmetals.

• Ionic Compounds – opposite sides of

periodic table (metal+nonmetal)

• Covalent Compounds – same side of

periodic table (nonmetal+nonmetal)

Diatomic Molecules

• These seven elements occur naturally as molecules

containing two atoms: all halogens + HON

– Hydrogen

– Oxygen

– Nitrogen

– Fluorine

– Chlorine

– Bromine

– Iodine

Types of Formulas

• Empirical formulas give the lowest whole-number

ratio of atoms of each element in a compound. (ex:

HO)

• Molecular formulas give the exact number of atoms

of each element in a compound. (ex: H2O2)

• If we know the molecular formula of a compound, we

can determine its empirical formula. The converse is

not true without more information!

Ions

• When an atom of a group of atoms loses or gains

electrons, it becomes an ion.

• Cations are formed when at least one electron is lost.

Monatomic cations are formed by metals. (+ charge)

• Anions are formed when at least one electron is gained.

Monatomic anions are formed by nonmetals, except the

noble gases. (--charge)

Charges on Main Group Elements

Cation Charges for Typical Main-Group Ions

1+ 2+ 3+

Charge = + Group #

( group 1A to 3A)

Charges on Main Group Elements

Cation Charges for Typical Main-Group Ions

1–2–3–charge = group # - 8

(for group 5A to 7A)

HW 2.3: What is the charge on the ions

formed from the following atoms ?

Na _______ group # ____charge

Ga ________ group # ____charge

N _________ group # ____charge

I ________ group # _____charge

As _______ group # ______charge

Common Anions – learn the polyatomic ions

in bold (formula, charge, name)

Table 2.5 Common Anionsa

a The ions we use most often are in boldface. Learn them first.

Writing Formulas (ionic compounds)

• Because compounds are electrically neutral, one can

determine the formula of a compound this way:

– The charge on the cation becomes the subscript on the

anion.

– The charge on the anion becomes the subscript on the

cation.

– If these subscripts are not in the lowest whole-number

ratio, divide them by the greatest common factor.

HW 2.4 : Write out the formula for the binary ionic

compound formed from the following elements.

Show work. (a) what are likely charges on all

atoms (b) write out correct ionic compound

formula including showing the subscript for how

many of each ion you need for a neutral formula

(zero charge formula) (c) give name

Ca & Cl

Rb & Se

Ga & O

Chemical Nomenclature

• The system of naming compounds is called chemical

nomenclature.

• We will learn how to name:

1) Ionic compounds

2) Binary Molecular Compounds (covalent)

3) Acids

summary for naming:▪ Ionic – metal nonmetal (if either is a polyatomic ion, use the

name of the polyatomic ion instead of the name of the element)

give element name for metal [If variable charge (most

transition metals & Sn, Pb & Tl) metal – use charge in parenthesis]

give element name for nonmetal – ending + ide

do NOT use number prefixes(Mg & S)

▪ Covalent – nonmetal-1 nonmetal-2 (same as ionic but use # prefix)

give element name for nonmetal-1

give element name for nonmetal-2 – ending + ide

use number prefixes (NO2)

(mono, di, tri, tetra, penta, hexa, hepta, octa, etc) ***

End1/26/21

Naming Chemical Compounds

Because nonmetals often combine

with one another in different

proportions to form different

compounds, numerical prefixes

are usually included in the names

of binary molecular compounds.

(covalent molecules)

Binary Molecular (covalent) Compounds

***

Ionic vs Covalent Naming

▪ Ca Cl2 vs P Cl3

▪ Ionic vs covalent

▪ (same naming but covalent use # prefix)

▪ calcium chlorine – ine + ide

calcium chloride

▪ phosphorus chlorine – ine + ide (use # prefix)

phosphorus trichloride

HW 2.5 : Naming Ionic Binary Compounds

1. Give the systematic name for each of the following compounds:

a. CoBr2

b. CaCl2

2. Given the following systematic names, write the formula for each compound:

a. Chromium(III) chloride

b. Gallium iodide

HW 2.5 : Naming Ionic Binary Compounds


a. CoBr2 cobalt (II) bromide

b. CaCl2 calcium chloride

2. Given the following systematic names, write the formula for each compound:

a. Chromium(III) chloride Cr Cl3b. Gallium iodide Ga I3

HW 2.6: Naming Compounds Containing Polyatomic Ions


a. Na2SO4

b. Mn(OH)2

2. Given the following systematic names, write the formula for each compound: (formula must equal zero charge overall)

a. Sodium carbonate

b. Sodium phosphate

HW 2.6: Naming Compounds ContainingPolyatomic Ions


a. Na2SO4 sodium sulfate

b. Mn(OH)2 manganese (II) hydroxide

2. Given the following systematic names, write the formula for each compound: (formula musts equal zero charge overall)

a. Sodium carbonate Na2 CO3

b. Sodium phosphate Na3 PO4

HW 2.7 : Naming Covalent Binary Compounds

1. Name each of the following compounds:

a. PCl5b. PCl3c. SO2

2. From the following systematic names, write the formula for each compound:

a. Sulfur hexafluoride

b. Sulfur trioxide


HW 2.7: Naming Covalent Binary Compounds

1. Name each of the following compounds:

a. PCl5 phosphorus pentachloride

b. PCl3 phosphorus trichloride

c. SO2 sulfur dioxide

2. From the following systematic names, write the formula for each compound:

a. Sulfur hexafluoride SF6

b. Sulfur trioxide SO3


Patterns in Oxyanion Nomenclature (1 of 3)

• When there are two oxyanions involving the same

element

– the one with fewer oxygens ends in –ite (-ous acid)

– the one with more oxygens ends in –ate (- ic acid)

▪−

2NO : nitrite; −

3NO : nitrate

▪−

3SO : sulfite; −

4SO : sulfate

Acid, Base Names (memorize * names & formulas)

*

*

*

*

*

*

*

*

Naming Acids

Naming Binary Acids (all halogens) (from last

page, binary acids of halogens below rule)

HCl hydrochloric acid

HBr hydrobromic acid

HF hydrofluoric acid (only weak halogen acid)

HI hydroiodic acid

Chemistry: The Central Science, 14e

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