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General Chemistry Study Guide

Chapter 2. Atoms, Molecules and Ions

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Yu Wang

1. Atoms

1.1. The Structure of Atoms

OpenStax 2.1 Early Ideas in Atomic Theory; 2.2 Evolution of Atomic Theory. Brown 2.1 The Atomic Theory of Matter; 2.2 The Discovery of Atomic Structure.

Dalton’s Atomic Theory (1808)

1. Elements are composed of extremely small particles called atoms.
2. All atoms of a given element are identical, having the same size, mass and chemical properties. The atoms of one element are different from the atoms of all other elements.
3. Compounds are composed of atoms of more than one element. In any compound, the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction.
4. A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction.

The Structure of Atoms

An atom is the smallest constituent unit of ordinary matter that has the properties of a chemical element. Every solid, liquid, gas, and plasma is composed of neutral or ionized atoms. Atoms are very small; typical sizes are around 100 picometers.

Every atom is composed of a nucleus and one or more electrons bound to the nucleus. The nucleus is made of one or more protons and typically a similar number of neutrons. Protons and neutrons are called nucleons. More than 99.94% of an atom's mass is in the nucleus. The protons have a positive electric charge, the electrons have a negative electric charge, and the neutrons have no electric charge. If the number of protons and electrons are equal, that atom is electrically neutral. If an atom has more or fewer electrons than protons, then it has an overall negative or positive charge, respectively, and it is called an ion.

Rutherford’s Model of the Atom

Compare the mass and charges of subatomic particles

Particle	Mass (g)	Charge (C)	Charge unit
Electron	$9.10938\times 10^{-28}$	$-1.6022\times 10^{-19}$	$-1$
Proton	$1.67262\times 10^{-24}$	$+1.6022\times 10^{-19}$	$+1$
Neutron	$1.67493\times 10^{-24}$	$0$	$0$

Radioactivity

(This section will be discussed further in Chapter 21 Nuclear Chemistry. Just have a brief idea of different types of radiations here.)

Radioactivity describes the spontaneous emission of particles and/or radiation. Consequently, any element that spontaneously emits radiation is said to be radioactive.

$\alpha$-radiation: Alpha radiation occurs when an atom undergoes radioactive decay, giving off a particle (called an alpha particle) consisting of two protons and two neutrons (essentially the nucleus of a helium-4 atom).

$\beta$-radiation: Beta radiation takes the form of either an electron or a positron (a particle with the size and mass of an electron, but with a positive charge) being emitted from an atom.

$\gamma$-radiation: Gamma rays are penetrating electromagnetic radiation of a kind arising from the radioactive decay of atomic nuclei. It consists of photons in the highest observed range of photon energy.

X-ray: X-radiation is a form of electromagnetic radiation. Most X-rays have a wavelength ranging from 0.01 to 10 nanometers, corresponding to frequencies in the range 30 petahertz to 30 exahertz and energies in the range 100 eV to 100 keV. X-Rays are longer-wavelength and (usually) lower energy than gamma radiation.

Important Experiments

(Just to have a brief idea how the subatomic structures were discovered.)

Measurement of the mass/charge ratio of electrons by J.J. Thomson

Cathode ray is a beam of electrons emitted from the cathode of a high-vacuum tube.

Cathode ray tube is a high-vacuum tube in which cathode rays produce a luminous image on a fluorescent screen, used chiefly in televisions and computer terminals. In some experiments, two electrically charged plates and a magnet were added to the outside of the cathode ray tube. The external electric field and magnetic field can change the trajectory of the cathode ray (electrons). The ratio of electric charge to the mass of an individual electron can be calculated based on electromagnetic theory.

Measurement of the mass of electrons by Millikan

The experiment entailed observing tiny electrically charged droplets of oil located between two parallel metal surfaces. when the gravitational force and external electric force are balanced, the charge on the oil droplet can be calculated. By repeating the experiment for many droplets, they confirmed that the charges were all small integer multiples of a certain base value, which was found to be $1.5924(17)\times10^{−19}\,\text{C}$, about $0.6\%$ difference from the currently accepted value of $1.602176487(40)\times10^{−19}\,\text{C}$.

Rutherford’s Experiment (This is important!)

Rutherford carried out a series of experiments using very thin foils of gold and other metals as targets for α particles from a radioactive source. He observed that the majority of particles penetrated the foil either undeflected or with only a slight deflection. In some instances, an α particle actually bounced back in the direction from which it had come! To explain the results of the α-scattering experiment, Rutherford devised a new model of atomic structure, suggesting that most of the atom must be empty space. This structure would allow most of the α particles to pass through the gold foil with little or no deflection. The atom's positive charges, Rutherford proposed, are all concentrated in the nucleus, a dense central core within the atom. Whenever an α particle came close to a nucleus in the scattering experiment, it experienced a large repulsive force and therefore a large deflection. Moreover, an α particle traveling directly toward a nucleus would experience an enormous repulsion that could completely reverse the direction of the moving particle.

Chadwick’s Experiment

It was known that hydrogen, the simplest atom, contains only one proton and that the helium atom contains two protons. Therefore, the ratio of the mass of a helium atom to that of a hydrogen atom should be 2:1. (Because electrons are much lighter than protons, their contribution can be ignored.) In reality, however, the ratio is 4:1. Rutherford and others postulated that there must be another type of subatomic particle in the atomic nucleus; the proof was provided by another English physicist, James Chadwick, in 1932. When Chadwick bombarded a thin sheet of beryllium with α particles, a very high energy radiation similar to γ rays was emitted by the metal. Later experiments showed that the rays actually consisted of electrically neutral particles having a mass slightly greater than that of protons. Chadwick named these particles neutrons.

Requirements

1. Understand Dalton’s Atomic Theory.
2. Understand Rutherford’s Experiment.
3. Understand the structure of atoms. Understand the correlations in the table above (Compare the mass and charges of subatomic particles).

1.2. Atomic Number, Mass Number, and Isotopes

OpenStax 2.3 Atomic Structure and Symbolism; 2.5 The Periodic Table. Brown 2.3 The Modern View of Atomic Structure; 2.5 The Periodic Table.

Atomic Number and Mass Number

The number of protons in the nucleus of each atom of an element is called the atomic number (Z).

$$\text{atomic number}=\text{number of protons}$$

The mass number (A) is the total number of neutrons and protons present in the nucleus of an atom of an element.

$$\text{mass number}=\text{number of protons}+\text{number of neutrons}$$

The accepted way to denote the atomic number and mass number of an atom of element X is as follows:

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Isotopes

Atoms that have the same atomic number but different mass numbers are called isotopes.

Periodic Table

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

The elements are arranged by atomic number (shown above the element symbol) in horizontal rows called periods and in vertical columns known as groups or families, according to similarities in their chemical properties.

The elements can be divided into three categories — metals, nonmetals, and metalloids. A metal is a good conductor of heat and electricity, whereas a nonmetal is usually a poor conductor of heat and electricity. A metalloid has properties that are intermediate between those of metals and nonmetals.

The Group 1A elements (Li, Na, K, Rb, Cs, and Fr) are called alkali metals, and the Group 2A elements (Be, Mg, Ca, Sr, Ba, and Ra) are called alkaline earth metals. Elements in Group 7A (F, Cl, Br, I, and At) are known as halogens, and those in Group 8A (He, Ne, Ar, Kr, Xe, and Rn) are called noble gases (or rare gases).

Requirements

1. Understand and remember the concepts.
2. Know the relation between atomic number and mass number.

2. Molecules and Ions

OpenStax 2.6 Molecular and Ionic Compounds. Brown 2.6 Molecules and Molecular Compounds; 2.7 Ions and Ionic Compounds.

2.1. General Concepts

Of all the elements, only the six noble gases in Group 8A of the periodic table (He, Ne, Ar, Kr, Xe, and Rn) exist in nature as single atoms. For this reason, they are called monatomic (meaning a single atom) gases. Most matter is composed of molecules or ions formed by atoms.

Molecules

A molecule is an aggregate of at least two atoms in a definite arrangement held together by chemical forces (also called chemical bonds).

The hydrogen molecule, symbolized as $\ce{H2}$, is called a diatomic molecule because it contains only two atoms. Molecules containing more than two atoms are called polyatomic molecules.

Organic Compounds

Organic chemistry is the branch of chemistry that deals with carbon compounds. The simplest type of organic compounds is the hydrocarbons, which contain only carbon and hydrogen atoms.

Example: methane $\ce{CH4}$

The chemistry of organic compounds is largely determined by the functional groups, which consist of one or a few atoms bonded in a specific way. For example, when an $\ce{H}$ atom in methane is replaced by a hydroxyl group ($\ce{—OH}$), an amino group ($\ce{—NH2}$), and a carboxyl group ($\ce{—COOH}$), the following molecules are generated: methanol ($\ce{CH3OH}$), methylamine ($\ce{CH3NH2}$), acetic acid ($\ce{CH3COOH}$).

Ions

An ion is an atom or a group of atoms that has a net positive or negative charge.

The loss of one or more electrons from a neutral atom results in a cation, an ion with a net positive charge. For example, a sodium atom ($\ce{Na}$) can readily lose an electron to become a sodium cation, which is represented by $\ce{Na+}$.

An anion is an ion whose net charge is negative due to an increase in the number of electrons. A chlorine atom ($\ce{Cl}$), for instance, can gain an electron to become the chloride ion $\ce{Cl^{−}}$.

Sodium chloride ($\ce{NaCl}$), ordinary table salt, is called an ionic compound because it is formed from cations and anions.

Ions like $\ce{Na+}$ and $\ce{Cl^{−}}$ are called monatomic ions because they contain only one atom. Figure 2.10 shows the charges of a number of monatomic ions. With very few exceptions, metals tend to form cations and nonmetals form anions.

Polyatomic ions such as $\ce{OH^{−}}$ (hydroxide ion), $\ce{CN^{−}}$ (cyanide ion), and $\ce{NH4+}$ (ammonium ion) are ions containing more than one atom.

Requirements

1. Understand and remember the concepts.

2.2. Chemical Formuas

OpenStax 2.4 Chemical Formulas; 2.6 Molecular and Ionic Compounds. Brown 2.6 Molecules and Molecular Compounds; 2.7 Ions and Ionic Compounds.

Chemists use chemical formulas to express the composition of molecules and ionic compounds in terms of chemical symbols.

Molecular Formulas

A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance.

An allotrope is one of two or more distinct forms of an element. Two allotropic forms of the element carbon—diamond and graphite—are dramatically different not only in properties but also in their relative cost.

Empirical Formulas

An empirical formula shows the simplest whole-number ratio of the atoms in a substance. For example, hydrogen peroxide $\ce{H2O2}$ has the empirical formula as $\ce{HO}$.

Molecular Models

Molecules are too small for us to observe directly. An effective means of visualizing them is by the use of molecular models.

A methane molecule can be represented as (a) a molecular formula, (b) a structural formula, (c) a ball-and-stick model, and (d) a space-filling model. Carbon and hydrogen atoms are represented by black and white spheres, respectively.

Formulas of Ionic Compounds

Ionic compounds consist of a combination of cations and anions.

• The formula is usually the same as the empirical formula.
• The sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero.

Example: Determine the formula of lead(IV) Oxide.
Answer:
The cation is $\ce{Pb^4+}$ and the oxygen anion is $\ce{O^{2−}}$. The following diagram helps us determine the subscripts for the compound formed by the cation and the anion:

The formula of the compound is $\ce{Pb2O4}$. We should reduce the subscripts to the lowest ratio, and get $\ce{PbO2}$. The sum of the charges is $1\times(+4) + 2(−2) = 0$. Thus, the formula for led(IV) oxide is $\ce{PbO2}$.

More examples see here.

Requirements

1. Understand and remember the concepts.

3. Naming Compounds

OpenStax 2.7 Chemical Nomenclature. Brown 2.8 Naming Inorganic Compounds.

3.1. Ionic Compouds

Often a metal and a nonmetal.

• Use the element name to represent the metal cation.
• Use the element name add "-ide" to represent the nonmetal anion.

For example, $\ce{NaCl}$ is named as sodium chloride.

For transition metals, indicate the charge on metal with Roman numberals. For example, $\ce{FeCl2}$ is named as iron(II) chloride, $\ce{FeCl3}$ is named as iron(III) chloride.

Memorize the bold ones in the following table:

Table $\mathbf{2.8}\mathbf{1}$: Common Polyatomic Ions and Their Names

Formula	Name of Ion	Formula	Name of Ion
NH4+	ammonium	HPO42−	hydrogen phosphate
CH3NH3+	methylammonium	H2PO4−	dihydrogen phosphate
OH−	hydroxide	ClO−	hypochlorite
O22−	peroxide	ClO2−	chlorite
CN−	cyanide	ClO3−	chlorate
SCN−	thiocyanate	ClO4−	perchlorate
NO2−	nitrite	MnO4−	permanganate
NO3−	nitrate	CrO42−	chromate
CO32−	carbonate	Cr2O72−	dichromate
HCO3−	hydrogen carbonate, or bicarbonate	C2O42−	oxalate
SO32−	sulfite	HCO2−	formate
SO42−	sulfate	CH3CO2−	acetate
HSO4−	hydrogen sulfate, or bisulfate	C6H5CO2−	benzoate
PO43−	phosphate

3.2. Molecular Compounds

Usually, molecular compounds are coumpounds with nonmetals or nonmetals + metalloids

• Common names may be used, such as water $\ce{H2O}$, ammonia $\ce{NH3}$, methane $\ce{CH4}$.
• Element furthest to the left in a period and closest to the bottom of a group on periodic table is placed first in formula.
• If more than one compound can be formed from the same elements, use prefixes to indicate number of each kind of atom, such as carbon monoxide $\ce{CO}$ and carbon dioxide $\ce{CO2}$.
• Last element name ends in -ide.

Greek prefixes:

Prefix	Meaning
mono-	1
di-	2
tri-	3
tetra-	4
penta-	5
hexa-	6
hepta-	7
octa-	8
nona-	9
deca-	10

3.3. Acids

An acid can be defined as a substance that yields hydrogen ions ($\ce{H+}$) when dissolved in water.

For example, $\ce{HCl}$ gas as a pure substance is named as hydrogen chloride. When $\ce{HCl}$ is dissolved in water, hydrochloric acid is formed giving $\ce{H3O+}$ and $\ce{Cl^{-}}$.

(Most complicated part!) An oxoacid is an acid that contains hydrogen, oxygen, and another element. The formulas of oxoacids are usually written with the H first, followed by the central element and then O. We use the following five common acids as our references in naming oxoacids:

Acid	Acid Name	Acid	Acid Name
$\ce{H2CO3}$	carbonic acid	$\ce{H3PO4}$	phosphoric acid
$\ce{HClO3}$	chloric acid	$\ce{H2SO4}$	sulfuric acid
$\ce{HNO3}$	nitric acid

Often two or more oxoacids have the same central atom but a different number of O atoms. Starting with our reference oxoacids, whose names all end with “-ic,” we use the following rules to name these compounds.

1. Addition of one O atom to the “-ic” acid: The acid is called “per… -ic” acid. Thus, adding an O atom to $\ce{HClO3}$ changes chloric acid to perchloric acid, $\ce{HClO4}$.
2. Removal of one O atom from the “-ic” acid: The acid is called “-ous” acid. Thus, nitric acid, $\ce{HNO3}$, becomes nitrous acid, $\ce{HNO2}$.
3. Removal of two O atoms from the “-ic” acid: The acid is called “hypo… -ous” acid. Thus, when $\ce{HBrO3}$ is converted to $\ce{HBrO}$, the acid is called hypobromous acid.

The rules for naming anions of oxoacids, called oxoanions, are

1. When all the H ions are removed from the “-ic” acid, the anion's name ends with “-ate.” For example, the anion derived from $\ce{H2CO3}$ is called carbonate.
2. When all the H ions are removed from the “-ous” acid, the anion's name ends with “-ite.” Thus, the anion derived from $\ce{HClO2}$ is called chlorite.
3. The names of anions in which one or more but not all of the hydrogen ions have been removed must indicate the number of H ions present. For example, consider the anions derived from phosphoric acid:

$\ce{H3PO4}$ : Phosphoric acid

$\ce{HPO4^{2-}}$ : Hydrogen phosphate

$\ce{H2PO4^{-}}$ : Dihydrogen phosphate

$\ce{PO4^{3-}}$ : Phosphate

Names of Oxoacids and Oxoanions That Contain Chlorine

Acid	Acid Name	Corresponding Anion	Anion Name
$\ce{HClO4}$	perchloric acid	$\ce{ClO4-}$	perchlorate
$\ce{HClO3}$	chloric acid	$\ce{ClO3-}$	chlorate
$\ce{HClO2}$	chlorous acid	$\ce{ClO2-}$	chlorite
$\ce{HClO}$	hypochlorous acid	$\ce{ClO-}$	hypochlorite

3.4 Bases

A base can be defined as a substance that yields hydroxide ions ($\ce{OH-}$) when dissolved in water, such as sodium hydroxide $\ce{NaOH}$.

3.5 Hydrates

Hydrates are compounds that have a specific number of water molecules attached to them, such as copper sulfate pentahydrate $\ce{CuSO4}\cdot\ce{5H2O}$.

Requirements

1. Understand and remember the concepts.

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