Post on 22-Oct-2021
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.
Postulates of Dalton’s Atomic Theory (2 of 4)
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)
Postulates of Dalton’s Atomic Theory (3 of 4)
3) Atoms are neither created nor destroyed in chemical
reactions. Atoms are just rearranged in chemical
reactions.
Postulates of Dalton’s Atomic Theory (4 of 4)
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
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)
Periodic Table (2 of 3)
• 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.
Periodic Table (3 of 3)
• 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
1. Give the systematic name for each of the following 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
1. Give the systematic name for each of the following compounds:
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
1. Give the systematic name for each of the following compounds:
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
Copyright © Cengage Learning. All rights reserved 51
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
Copyright © Cengage Learning. All rights reserved 52
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