chapter 2 atoms, molecules, and ions. the vocabulary of chemistry is specialized. it includes such...

Chapter 2

Atoms, Molecules, and Ions

Atoms, Molecules, and Ions

• The vocabulary of chemistry is specialized.

• It includes such terms as atom, ion, molecule, isotope, acid,

base, salt, and saturated hydrocarbon.

• Chemistry also involves symbolic notation.

• Atoms of elements are represented by symbols, such as H,

C, N, O, Na, and Cl. These symbols are the alphabet of

chemistry.

• To represent compounds, symbols are combined into

chemical formulas, such as NaCl, and formulas are the words

of chemistry.

Gen. Chem. Chapter2 3

Molecule models

S 8 C 4


Molecular compounds


Some moleculesH2O2 CH3CH2Cl P4O10

CH3-CH(OH)-CH3 HCOOH


Molecule models


New laws are formulated making use of other laws.

BRICKS ON THE WALL!!!!

• Is the mass of the burned match pictured in the photograph the same as, less than, or

more than the mass of the burning match?

The match loses mass in the

reaction. This kind of loss of mass

to carbon dioxide during burning

greatly confused early chemists

who did not measure the mass of

the gases produced.

The rule of conservation of matter

is credited to Lavoisier.


Gen. Chem. Chapter2 9Antoine Lavoisier (1743 – 1794)


Another of the 18th century chemists to emphasize the importance of

experiments was Antoine Lavoisier (1743-1794)

Lavoisier studying human respiration. His wife, Marie Anne Paulze,

seated at the table on the right, records the experiment. Lavoisier

was the first great theoretical chemist. Marie was an accomplished

artist and made this drawing herself.


2.1 Laws of Chemical Combination• Data obtained by experiment can often be summarized into

laws. Scientific theories are then formulated to explain these laws.

Lavoisier: The Law of Conservation of Mass

• Modern chemistry dates from the eighteenth century, when

scientists began to make quantitative observations.

• Antoine Lavoisier made mass measurements precise to about

0.0001 g and established that total mass does not change

during a chemical reaction.


• For example, Lavoisier heated the red oxide of mercury,

causing it to decompose (break down) to two new products:

mercury metal and oxygen gas.

• By measuring carefully, he found that the total mass of the

products was exactly the same as the mass of the mercury

oxide he started with.

• Lavoisier summarized his findings through the

law of conservation of mass.

The total mass remains constant during a chemical

reaction.


Proust: The Law of Definite Proportions

• By the end of the eighteenth century, Lavoisier and other scientists

had succeeded in decomposing many compounds into the

elements that form them.

• One of these scientists, Joseph Proust (1754–1826), did careful

quantitative studies by which he established the

Law of Constant Composition

• (also called the Law of Definite Proportions).

All samples of a compound have the same composition; that is, all

samples have the same proportions, by mass, of the elements

present in the compound.


Joseph Proust (1754 – 1826)


• Proust based this law in part on his studies of a substance that we

now call

• basic copper carbonate ;

• all samples of this compound have the same composition:

• 57.48% copper,

• 5.43% carbon,

• 0.91% hydrogen, and

• 36.18% oxygen by mass.

Gen. Chem. Chapter2

The law of definite proportions : Regardless of its source, basic copper carbonate Cu2(CO3)(OH)2, always has the same

composition. 16

A substance known as basic copper carbonate occurs in nature

as the mineral malachite forms as a patina on copper roofs and bronze statues

can be synthesized in the laboratory


•Chemistry: Cu2(CO3)(OH)2, Copper Carbonate Hydroxide. •Class: Carbonate

The Mineral Malachite


• A compound not only has a constant, or fixed, composition; it

also has fixed properties.

• For example, under normal atmospheric pressure, pure water

always freezes at 0 °C and boils at 100 °C.


• The physical and chemical properties of chemical substances—both

elements and compounds—depend on their composition.

• A twentieth-century modification of this law, based on the work of

Albert Einstein, allows for the possibility of converting mass to energy

(and vice versa).

• This need not concern us, however, because in chemical rections the

amounts of mass converted to energy are too small to detect.


2.2 John Dalton and

the Atomic Theory of Matter

In 1803, John Dalton proposed a theory to

explain

the laws of conservation of mass and constant

composition.

As he developed what would become known as his atomic theory,

Dalton found evidence of a scientific law describing the composition of

matter.

In some cases, atoms of the same two elements are able to combine to

form two or more different compounds, and the compositions of the

different compounds are related.


John Dalton


• Let's illustrate Dalton's reasoning with a simple example: two compounds

composed of the elements carbon and oxygen.

• In carbon dioxide (the familiar gas produced in respiration and in the

burning of fuels) the two elements combine in the ratio of 8.0 g of oxygen to

3.0 g of carbon.

• In carbon monoxide (the poisonous gas formed when a fuel is burned in

limited air) the elements combine in the ratio of 4.0 g of oxygen to 3.0 g of

carbon.

• Based on evidence such as this, Dalton formulated

The Law of Multiple Proportions• which we can state in the following way:

In two or more compounds of the same two elements, the masses of

one element that combine with a fixed mass of the second element

are in the ratio of small whole numbers.


Law of Multiple Proportions

Ratio of oxygen-

to-carbon in CO2

is exactly twice the

ratio in CO

ANIMATION 1: MULTIPLE PROPORTIONS



Law of Multiple Proportions

• Four different oxides of nitrogen can be formed,

by combining 28 g of nitrogen with:

• 16 g oxygen, forming Compound I

• 48 g oxygen, forming Compound II

• 64 g oxygen, forming Compound III

• 80 g oxygen, forming Compound IV

• Compounds I–IV are N2O, N2O3, N2O4, N2O5

What is the ratio 16:48:64:80

expressed as small whole numbers?1 : 3 : 4 : 5


• Dalton's Atomic Theory

• The atomic model Dalton developed to explain the laws of chemical

combination is based on four ideas:

- All matter is composed of extremely small, indivisible particles called

atoms.

- All atoms of a given element are alike in mass and other properties, but

the atoms of one element differ from the atoms of every other element.

- Compounds are formed when atoms of different elements unite in fixed

proportions. That is, one atom of A to one of B in AB, two atoms of A to

one of B in A2B and so on.

- A chemical reaction involves a rearrangement of atoms to produce new

compounds. No atoms are created, destroyed, or broken apart in a

chemical reaction.


2.3 The Divisible Atom

• However, like most other scientific theories, Dalton's model

eventually had to be modified in light of later discoveries.

• Toward the end of the nineteenth century, certain experiments

began to reveal that atoms are made up of smaller parts.

• Of the dozens of subatomic (smaller than atomic) particles now

known, three are of special importance in the study of

chemistry.

• They are the proton, neutron, and electron.


Subatomic Particles

• The mass and charge of subatomic particles are so small that they are

conveniently expressed in relative units.

• The proton has a relative mass of 1. The proton also carries one fundamental

unit of positive electric charge, denoted 1+ .

• The neutron, as its name implies, is an electrically neutral particle; it has no

charge. Although its mass is slightly greater than that of the proton, for many

purposes we can consider the neutron also to have a relative mass of 1.

• The third particle, the electron, has a mass that is 1/1836 of the mass of a

proton. An electron has the same quantity of charge as a proton, but it is a

negative charge, denoted as 1- .

• Protons, neutrons, and electrons are fundamental particles.

• This means that all protons are alike, all neutrons are alike, and all electrons

are alike in whatever element they are found.


Subatomic Particles

• Protons and neutrons are densely packed into

a tiny, positively charged core of the atom

known as the nucleus.

• The extremely lightweight electrons are widely

dispersed around the nucleus.

• An atom as a whole is electrically neutral: It has

no net charge because the negative and positive

charges balance each other.Atom like Football Field


Atomic Diameter 10-8 cm

Nuclear Structure

= 1 Å

Nuclear diameter 10-13 cm


Subatomic Particles

•Although the numbers differ from one element

to another, the atoms of every element have

equal numbers of electrons and protons.


• Dalton believed that the identity of an element is determined by

the mass of one of its atoms.

• We now know that it is not the mass of an atom but rather the

number of protons in the nucleus that determines the kind of

atom and therefore the identity of an element.

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

of an atom of a given element, and it is this number of protons

that defines the element.

• Every atom having two protons in its nucleus has Z = 2 and is

an atom of helium.

• Every atom with 92 protons has Z = 92 and is an atom of

uranium.


• Atoms that have the same number of protons but different

numbers of neutrons are called isotopes.

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

of a given atom of a given element.

• The mass number (A) is an integral number that is the sum of

the numbers of protons and neutrons in an atom.

• The number of neutrons = A – Z.

Isotopes


Isotopes

• For example, there are three isotopes of hydrogen.

• The most abundant isotope, occasionally called protium, has a

single proton and no neutrons in its nucleus.

• About one in every 6700 hydrogen atoms, however, has a

neutron as well as a proton. Because the mass of a neutron is

essentially the same as the mass of a proton, the mass of this

hydrogen isotope, called deuterium, is about twice that of

protium.

• A third, very rare isotope of hydrogen, called tritium, has two

neutrons and one proton in the nucleus. A tritium atom has

about three times the mass of a protium atom.


Isotopes of Hydrogen

Atomic number (Z) = number of protons

Protium has 1 proton, 0 neutrons : Z = 1

Deuterium has 1 proton, 1 neutron : Z = 1

Tritium has 1 proton, 2 neutrons : Z = 1


• Some elements have only one naturally occurring isotope; these

include fluorine-19, sodium-23, and phosphorus-31.

• Most elements, however, have two or more isotopic forms.

• Tin has the greatest number of naturally occurring isotopes: 10.

• The naturally occurring isotopes are always found in certain

precise proportions. In natural sources of chlorine, for example,

75.77% of the atoms are chlorine-35 and 24.23% are chlorine-37.

• Chemical symbols for isotopes are commonly written in the form

with A being the mass number and Z the atomic number of

the element E.

• For the two naturally occurring isotopes of chlorine, we can write

and indicating mass numbers of 35 and 37

for chlorine-35 and chlorine-37, respectively.

A EZ

35 Cl1737 Cl17


Chemical Symbol

To represent a particular atom we use the symbolism:

A= mass number Z = atomic number

To represent an ion


• How many protons are in chlorine-35?• How many protons are in chlorine-37?• How many neutrons are in chlorine-37?

Atoms can be represented using the element’s symbol and the mass number (A) and atomic number (Z):

A EZ 37 Cl1735 Cl17

Isotopes


Other Examples of Isotopes

Carbon-14 Z = 6 so 8 neutrons

The number of neutrons = A – Z

Chlorine-35 Z = 17 so 18 neutrons

Uranium-234 Z = 92 so 142 neutrons

Carbon-12 Z = 6 so 6 neutronsCarbon-13 Z = 6 so 7 neutrons

Chlorine-37 Z = 17 so 20 neutrons

Uranium-235 Z = 92 so 143 neutronsUranium-238 Z = 92 so 146 neutrons


C, H, O Analyser


• By international agreement, the current atomic mass standard is

the pure isotope carbon-12, which is assigned a mass of exactly

12 atomic mass units (12 u).

• Based on this standard, we can define an atomic mass unit,

(abbreviated amu and having the unit u)

as exactly one-twelfth the mass of a carbon-12 atom.

• In more familiar units of mass, 1 u = 1.66054 x 10-24 g.

• Because we are interested in the masses of atoms and not their weights,

we will now shift to the term atomic mass in place of the older atomic

weight, except in a few cases of historical interest.

2.4 Atomic Masses


• The atomic mass of an element is defined as the weighted average of the

masses of the naturally occurring isotopes of that element.

• Consider carbon, which consists of a mixture of two naturally occurring

isotopes.

• The much more abundant isotope is carbon-12. The other is carbon-13,

with a mass of 13.00335 u. Both isotopes are present in substances

containing carbon atoms, and in proportions that generally do not vary

from one carbon-containing substance to another.

• To describe the atomic mass of carbon, then, we need to use an average

value, but not the simple average (12 + 13) / 2 = 12.5. Because carbon-12 is

much more abundant than carbon-13, the average we seek lies much

closer to the mass of carbon-12 than to that of carbon-13. We say that it is

"weighted" toward the mass of carbon-12.


• To calculate the atomic mass of an element, we need two quantities—

• (1) the atomic masses of the isotopes of the element and

• (2) the naturally occurring fractional abundances of the isotopes.

• Let us explain how these quantities are obtained experimentally

• To illustrate, let's return to the atomic mass of carbon.


• The percentage abundances and fractional abundances of the carbon isotopes are as follows:

To obtain a weighted average atomic mass, we calculate the contribution of each isotope to the weighted average from the relationship

The term atomic weight is still widely used, however, as by the Commission on Atomic Weights of the International Union of Pure and Applied Chemistry (IUPAC).


The contribution of each isotope to the weighted average atomic mass is given by Equation.

SolutionThe contributions are

Contribution of carbon-12 = 0.98892 x 12.00000 u = 11.867 u

Contribution of carbon-13 = 0.01108 x 13.00335 u = 0.1441 u

The weighted average atomic mass is the sum of the two contributions.

Atomic mass of carbon = 11.867 u + 0.1441 u = 12.011 u

Use the data cited below to determine the weighted average atomic mass of carbon.


Carbon – 14 ( C-14 Dating )


2.5 The Periodic Table: Elements Organized

• In the nineteenth century, chemists discovered dozens of new

elements.

• By 1830, 55 elements were recognized, but there was no

apparent pattern in their properties. Chemists badly needed a

way to organize the growing collection of chemical data.

• One way to do this was to arrange the elements in a manner

that would establish categories of elements having similar

physical and chemical properties.

• Dimitri Mendeleev published the first successful arrangement,

called a periodic table, in 1869.

• In its modern form, the periodic table organizes a vast array of

chemical knowledge.


Dimitri Mendeleev


Mendeleev's Periodic Table

Mendeleev arranged the elements in order of

increasing atomic weight, from left to right in rows

and from top to bottom in columns (or groups).

In this arrangement, elements that most closely

resemble one another in physical and chemical

properties tend to fall in the same vertical group.

This group similarity repeats periodically, hence the name periodic table.

• There would be no exceptions to the principle that all the elements in a

group display similar properties, therefore Mendeleev placed some

elements out of order, that is, not in the strict order of increasing atomic

weight.

For example, he correctly

placed tellurium (atomic weight

127.6) before iodine (atomic

weight 126.9) so that tellurium

would be in the same column

as the similar elements sulfur

and selenium.



• When Mendeleev placed elements with similar properties in

the same vertical group, a few gaps were created in his table.

Instead of seeing these gaps as defects, he boldly predicted

the existence of undiscovered elements to fill the gaps.

• Furthermore, because the table was based on patterns of

properties, he was able to predict some properties of the

missing elements.


Mendeleev’s TableDimitri Mendeleev created this, the original periodic table.

Stow

e's table S

piral form

Trian

gular form

Fold

ed T

able P

eriodic T

able

Mendeleevs’s early table was published in 1872.

• Using his table he was able to correct several values of atomic

masses.

• Because of its obvious usefullness his periodic table was almost

universally adopted, and it remains one of the most valuable tools at the

chemist’s use.

• The only fundemantal difference between todays table and that of his

is that in the current table the elements are ordered by atomic number

rather than by atomic mass.

Mendeleev’s Table


• For example, he left a blank

space for an undiscovered

element that he called "eka-

silicon" and used its location

between silicon and tin to

predict an atomic weight of 72

and other properties.


• Table 2.2 shows just how accurate his predictions were, when compared

to the properties of the actual element germanium discovered 15 years

later.

• The predictive nature of Mendeleev's periodic table led to its wide

acceptance as a tremendous scientific accomplishment.


The Periodic tableAlkali Metals

Alkaline Earths

Transition Metals

Halogens

Noble Gases

Lanthanides and Actinides

Main Group

Main Group

(2) PERIODIC

PROPERTIES

ANIMATION



Empirical and Molecular Formulas

Empirical formula: the simplest whole number ratio of elements in a compound

Example: Molecular formula of glucose – C6H12O6

The elemental ratio C:H:O is 1:2:1, so the empirical formula is CH2O


Molecular compounds


Introduction to Compounds

A molecule is a group of two or more atoms held together by covalent bonds.

A chemical formula is a symbolic representation of the composition of a compound in terms of its constituent elements.


Structural Formulas

Shows how atoms are attached to one another.


Molecular Compounds

Ball-and-stick model vs. Space-filling model


Binary Molecular Compounds

Compounds that are typically comprised of two

nonmetallic elements:

e.g., CO, NO, HF

Molecular formulas are usually written with the

more “metallic” first – “metallic” means farther

left in the period and lower in the group

e.g., NaCl, KBr


Formulas and Subscripts

Subscripts are used when a given atom is used more than once

e.g., H2O, CO2, N2O, HF, B2O3

The presence of subscripts is reflected in the names of compounds


Names of Binary Compounds

The compound name consists of two words, one for each element in the compound

Consider the compounds CO and CO2

Name the element that appears first in the formula: CARBON

The second element has an altered name: retain the stem of the element name and replace the ending by -ide

OXYGEN OXIDE

However, both compounds cannot be carbon oxide


Names of Binary CompoundsThe names are further modified by adding prefixes

to denote the numbers of atoms:

CO (Carbon Mon-oxide), CO2 (Carbon Di-oxide)


2.7 Ions and Ionic Compounds

In an isolated atom, the number of protons equals the number of

electrons, and the atom is therefore electrically neutral.

In some chemical reactions, however, an individual atom or a

group of bonded atoms may lose or gain one or more electrons,

thereby acquiring a net electric charge and becoming an ion.

Ions are formed only through the loss or gain of electrons; there

is no change in the number of protons in the nucleus of the

atom(s). If electrons are lost, there are more protons than

electrons in the resulting ion and so it has a positive charge.

If electrons are gained, there are more electrons than protons in

the resulting ion and so it has a negative charge.


Ions and Ionic Compounds

Atoms that gain or lose electrons are called ions

Positive ions: CATIONS Negative ions: ANIONS

Atoms that lose electrons form cations

Na Na+ + e–

Atoms that gain electrons form anions

Cl + e– Cl–


Monatomic Ions

Group A metals usually lose the number of electrons equal to

their Group number.

Nonmetal atoms usually gain electrons and have a charge equal

to their Group number minus eight.

The periodic table cannot be used to determine the charge on

Group B metals.

For naming, Group B metals capable of multiple charges have

the corresponding Roman numeral in parentheses added after

the element name.


Common Monatomic IonsFIGURE 2.10 Symbols and periodic table locations of some monatomic ions

Three general observations can be made of these data:

(1) Aluminum and the metals of groups 1A and 2A form just one cation, which carries a positive charge equal in magnitude to the A-group number.

(2) Most of the metals of the B groups form two or more cations of different charges, though in some cases only one of these cations is commonly encountered.

(3) The nonmetals of groups 7A and 6A, along with nitrogen and phosphorus of group 5A, form anions that have a charge equal to the group number minus 8.


Formulas and Names forIonic Compounds

Ionic compounds form when oppositely charged ions are attracted to each other NaCl

Resulting compound is electrically neutral

Na+ Cl–

(+1) + (–1) = 0

Ionic compound names use the cation name followed by the anion name

Sodium chloride


Now consider the formula for aluminum oxide.

We cannot simply combine one Al3+ and one O2- , because

this would produce a formula unit with a net charge of 1+.

The combination of two Al3+ ions and three O2- ions, though,

is an electrically neutral formula unit:

2 (3+) + 3 (2-) = + 6 - 6 = 0

Therefore, the formula for aluminum oxide is Al2O3 .


Polyatomic Ions

Polyatomic ions are charged groups of covalently bonded atoms


Hydrates

A hydrate is an ionic compound in which the formula unit includes a fixed number of water molecules associated with cations and anions

Examples:

BaCl2 . 2 H2O

CuSO4 . 5 H2O


Acids

• Taste sour

• Turn blue litmus paper red

• React with metals to form

hydrogen gas

• Neutralize a base

IntroToAcids


Bases

• Taste bitter

• Turn red litmus paper blue

• Feel slippery on skin

• Neutralize an acid

IntroToBases


Arrhenius Concepts

Acids are compounds that ionize in

water to form a solution of H+ ions

and anions

Bases are compounds that ionize in

water to form solutions of OH– and

cations


Arrhenius Concepts

HCl + NaOH “Salt” + Water

Acid Base Na

cation

HOHCl

/anion

Acids and bases react to form a salt and water

= neutralization

Animation (3) & (4) : Acids, Bases



Formulas and Names for Acids

Binary acids start with

hydro and end with “ic”

plus the word acid

Ternary acids simply use

the polyatomic anion

name with “ate”

changing to “ic” plus the

word acid


Formulas and Names for Bases

Arrhenius bases always have hydroxide ions

The name follows ionic compound convention

e.g., NaOH – sodium hydroxide

Molecular bases form OH– after reacting with water

NH3 + HOH NH4OH

Ammonia ammonium hydroxide


Formulas and Names for Salts

Binary salts use the

“ide” ending on the

anion name

e.g., sodium chloride

Polyatomic salts use

“ate” ending on the

anion name

e.g., sodium sulfate


Organic Compounds

• Organic Chemistry is the study of carbon and its

compounds

• Carbon compounds containing one or more of the

elements H, O, N, or S are especially common

• Most organic compounds are molecular compounds

• Can exist as acids, bases, and salts

• Compounds have systematic names AND common

names


Representations of Molecules

Condensed Structural Formula CH3CH2CH3

Structural Formula

Ball and Stick


Saturated Hydrocarbons

Hydrocarbons have only hydrogen and carbon atoms

Saturated hydrocarbon: has

the maximum number of

hydrogen

atoms possible for each carbon

atomAlkanes are saturated hydrocarbons

Methane (CH4) is the first molecule in the

alkane series


Prefixes for Number of Carbon

Used for simple organic molecules

Combined with alkane ending “ane”

e.g., propane is a 3-carbon alkane


Isomers

Compounds with the same molecular formula but

different structural formulas


FIGURE 2.14 Butane and isobutaneBall-and-stick and structural models illustrate the

difference between the two isomers of butane.


Cyclic Alkanes

Alkane compounds that have carbons arranged in a ring structure are called cycloalkanes.

use the prefix cyclo-

methylcyclopropane

cyclohexane


FIGURE 2.15 Cyclohexane

The molecular and line-angle formulas for cyclohexane indicate the bonding of atoms

within the molecule. The ball-and-stick model further indicates that all of the carbon atoms

of the cyclohexane molecule do not lie in the same plane. Rather, it can assume several

different arrangements or conformations. The most stable arrangement, shown here, is

called the chair conformation because the six carbon atoms outline a structure that

somewhat resembles a reclining chair.


Functional Groups

• Alcohols• Ethers• Carboxylic Acids• Esters• Amines

Specific groupings of atoms attached to a carbon chain that give the compound unique properties

Most-common functional groups include:


AlcoholsAlcohols are molecules that contain a hydroxyl group (OH)


Alcohols


Ethers

Ethers are molecules in which two alkane groups (R-) are attached to a central oxygen atom

The general formula is R-O-R´

R and R´ may be the same or different groups

CH3CH2OCH2CH3

CH3CH2OCH2CH2CH3


Carboxylic Acids

Carboxylic acids are alkanes that also contain a carboxyl group and are weak acids

HCOO– + H+

Acts like an Arrhenius acid, loses a hydrogen ion


Esters

Esters are molecules in which two alkanes are attached to each side of a carboxyl group (R’-COO-R)


NH2(CH2)4NH2

Amines

Amines are molecules in which alkanes and hydrogen(s) are attached to a central nitrogen

Amines are weak bases


Summary of Concepts

• The basic laws of chemical combination are the laws of

conservation of mass, constant composition, and

multiple proportions.

• The three main subatomic particles are the protons,

neutrons, and electrons.

• Atoms with the same number of protons but different

numbers of neutrons are called isotopes.

• A chemical formula indicates the relative numbers of

atoms of each type in a compound.


Summary (cont.)

• The periodic table is an arrangement of the elements

by atomic number that places elements with similar

properties into the same vertical group.

• Ions are formed by the gain or loss of electrons.

Positive ions are cations and negative ions are anions.

• Many compounds are classified as either acids (H+),

bases (OH–), or salts (neutralization of acid and base).

• Organic compounds are based on the element carbon.

• Functional groups confer distinctive properties on an

organic molecule.


Arrhenius Concepts

HCl + NaOH “Salt” + Water

Acid Base Na

cation

HOHCl

/anion

Acids and bases react to form a salt and water

= neutralization

chapter 2 atoms, molecules, and ions. the vocabulary of chemistry is specialized. it includes such...

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