Chemistry · Chapter 2

Atoms, Molecules and Ions

Everything around you is built from a short list of parts. Protons, neutrons and electrons pack into atoms. Atoms join into molecules, or trade charge and become ions. From there they stack into every material you have ever touched. This chapter hands you the parts and the rules. You will count protons to name an element. You will mix isotopes to get an atomic mass. You will read the periodic table like a map. You will balance charge to write a formula. Then you will pick up the naming system chemists share. By the end, a label like Fe(NO₃)₃ will read like a sentence.

By the end of this chapter you can

Name the three subatomic particles and give the charge and mass of each.
Use atomic number and mass number to identify any atom or ion.
Work out an average atomic mass from isotope masses and abundances.
Find groups, periods, metals, nonmetals and metalloids on the table.
Predict ion charges and write neutral ionic formulas.
Tell empirical, molecular and structural formulas apart.
Name ionic, molecular and acid compounds from their formulas.

Section 1Atomic theory and the subatomic particles

A very old idea, finally tested

People argued about matter for a long time. Could you cut a lump of gold in half forever? Or do you reach a smallest piece that will not split? For centuries nobody could test the question. Careful weighing finally settled it.

In 1808 John Dalton pulled the evidence into one theory. An atom is the smallest piece of an element that still behaves like that element. All atoms of one element act the same way in a reaction. Atoms of different elements differ. Atoms combine in fixed, whole-number ratios to make compounds. And a chemical reaction only rearranges atoms. It never creates them and never destroys them.

Two mass laws backed him up. The law of definite proportions says one compound always holds the same mass ratio of elements. Pure water is 11.19% hydrogen by mass, whether it came from a river or a flask. The law of multiple proportions compares two compounds made from the same pair of elements. In carbon monoxide, 1.00 g of carbon is joined to 1.33 g of oxygen. In carbon dioxide, 1.00 g of carbon is joined to 2.66 g of oxygen. The ratio of 2.66 to 1.33 is a clean 2 to 1. Tidy whole numbers like that point to whole particles.

Three particles, and where they sit

Dalton thought the atom had no parts. He was wrong, and the next hundred years showed why. J. J. Thomson ran a current through a thin gas and watched a beam bend toward the positive plate. He had found the electron, a very light particle with a negative charge. Ernest Rutherford then fired tiny positive bullets at thin gold foil. Most flew straight through. A few bounced almost straight back. That only happens if the positive charge is crammed into a tiny core. He named the core the nucleus. The positive particle inside it is the proton. In 1932 James Chadwick found the neutron, a nuclear particle with no charge at all.

Masses here are quoted in the atomic mass unit (u). One unit is exactly 1/12 the mass of a carbon-12 atom, or 1.6605 × 10^-24 g. A proton weighs 1.0073 u. A neutron weighs 1.0087 u. An electron weighs only 0.00055 u, which is about 1/1836 of a proton. So the nucleus carries almost all of the mass.

Size runs the other way. A whole atom is roughly 10^-10 m across. Its nucleus is roughly 10^-15 m across. The atom is about 100,000 times wider than its own core. Blow the nucleus up to the size of a marble and the atom would stretch across a sports field. Almost all of an atom is empty space with a few electrons moving through it.

Inside an atom. The nucleus holds the protons and neutrons and nearly all of the mass. Electrons are drawn here in simple shells. The symbol carries the mass number on top and the atomic number below.

Counting protons names the element

The atomic number (Z) is the number of protons in the nucleus. It is the atom's identity badge. Six protons means carbon, every single time. The mass number (A) is the protons plus the neutrons added together. Chemists write both on the symbol, with A above and Z below. A neutral atom holds as many electrons as protons. Change the electron count and the atom carries a charge instead.

Particle	Where it sits	Charge	Mass (u)	Mass vs a proton
Proton	In the nucleus	1+	1.0073	1
Neutron	In the nucleus	0	1.0087	about 1
Electron	Around the nucleus	1−	0.00055	about 1/1836

Worked example · Reading an atom from its particle counts

A particle holds 26 protons, 30 neutrons and 24 electrons. What is it?

Step 1 · Use the protons to name itZ = 26, so this is iron, symbol Fe. The protons alone decide this.

Step 2 · Add for the mass numberA = 26 protons + 30 neutrons = 56.

Step 3 · Compare protons with electronscharge = 26 − 24 = 2+. Two electrons are missing, so two positive charges are left over.

Step 4 · Write it⁵⁶Fe²⁺, read as iron-56 with a 2+ charge. It is an ion, not a neutral atom.

CheckProtons still equal 26, so it is still iron. Only the charge and the neutron count changed.

Key points

Atoms combine in whole-number ratios, which is what the two mass laws showed.
Protons are 1+, neutrons are neutral, and both sit in the nucleus. Electrons are 1− and sit outside.
The nucleus holds nearly all the mass in nearly none of the volume.
Z counts protons and names the element. A counts protons plus neutrons.
Protons minus electrons gives the charge.

Interactive Add protons, neutrons and electrons and see what you have made. Open full ↑

Section 2Isotopes and atomic mass

Same element, different weight

Dalton said every atom of an element has the same mass. That part needed a repair. Atoms of one element can hold different numbers of neutrons. Those versions are isotopes. They share an atomic number and differ in mass number. Hydrogen shows it most clearly. Almost every hydrogen atom is one proton with no neutron at all. Roughly one in 8,700 carries a neutron, and that form is called deuterium. A third form, tritium, carries two neutrons and breaks down over time.

Isotopes of one element react the same way. Chemistry is run by electrons, and isotopes have matching electron counts. Mass does shift a few physical traits. Heavy water, D₂O, freezes at 3.8 °C instead of 0 °C. It is also about 11% denser than ordinary water.

A nuclide is one specific build of a nucleus, such as carbon-13 or uranium-235. Naming a nuclide means fixing both numbers, not just the element. Some nuclides are stable and last forever. Others come apart and release radiation.

Top: three isotopes of hydrogen, same one proton, different neutron counts. Bottom: bar height is how common each chlorine isotope is. The average sits where the two bars would balance, so it lands near the common lighter one.

Weighing a mixture

A mass spectrometer sorts atoms by mass. It knocks an electron off each atom, speeds the ions up, then bends their path with a magnet. Heavy ions bend less than light ones. A detector counts how many land at each mass. Out comes the mass of every isotope and its percent abundance, the share of atoms that are that isotope.

The number printed under each symbol on the periodic table is the average atomic mass. It is a weighted average, not a plain one. Multiply each isotope mass by its share, then add the results:

average atomic mass = Σ (fraction × isotope mass)

Use the fraction, so divide each percent by 100 first. The answer always leans toward the common isotope. Chlorine averages 35.45 u, far nearer 35 than 37, because the lighter form is about three times as common.

Worked example · The average atomic mass of chlorine

Chlorine is 75.76% chlorine-35 at 34.96885 u and 24.24% chlorine-37 at 36.96590 u. Find its average atomic mass.

Step 1 · Turn percents into fractions75.76% → 0.7576 and 24.24% → 0.2424. They add to 1.0000, so nothing is missing.

Step 2 · Weight each mass0.7576 × 34.96885 u = 26.4924 u.
0.2424 × 36.96590 u = 8.9605 u.

Step 3 · Add the pieces26.4924 u + 8.9605 u = 35.4529 u, which rounds to 35.45 u.

CheckThe answer must fall between 34.97 and 36.97, and it does. It sits close to the lighter isotope, which is the common one. That matches the table value of 35.45 u.

An atom weighs less than its parts. Add up the pieces of a carbon-12 atom and you get 6(1.0073) + 6(1.0087) + 6(0.00055) = 12.0993 u. Yet a carbon-12 atom is exactly 12 u, because the unit was built from it. The missing 0.0993 u was released as energy when the nucleus formed. Chemists call the gap the mass defect.

Mass number is a count, atomic mass is a mass

Keep these two apart. A mass number is a whole number and nothing else. It counts protons plus neutrons. An atomic mass is a measured mass in u, and it usually arrives with decimals. Carbon-13 has mass number 13 and an atomic mass of 13.00335 u. Only carbon-12 has both numbers equal, and that is on purpose.

Key points

Isotopes of one element have the same Z and different A, so they differ only in neutrons.
Isotopes act alike chemically, because they hold the same electrons.
Percent abundance is the share of atoms of that isotope in a natural sample.
Average atomic mass = sum of (fraction × isotope mass), so it always leans toward the common isotope.
Mass number counts particles. Atomic mass measures mass in u.

Interactive Mix isotopes to their natural abundance and get the average atomic mass. Open full ↑

Section 3The periodic table

A map built out of behavior

By the 1860s chemists knew dozens of elements and their masses. Dmitri Mendeleev lined them up in order of mass and spotted a repeat. Properties kept coming back around. So he broke the line into rows and stacked the look-alikes into columns. He even left holes for elements nobody had found, and guessed what they would be like. Gallium turned up in 1875 and matched his guess closely. That is when the table stopped looking like a filing trick and started looking like a law.

Today the elements are ordered by atomic number instead of mass. That fixed a few seats. Tellurium weighs more than iodine, yet tellurium comes first, because it has fewer protons. Order by protons and every look-alike lands in the right column.

Rows, columns and the long drawers

A row is a period, numbered 1 through 7. A column is a group, numbered 1 through 18. Elements in one group behave alike because they hold the same number of outer electrons. Four groups have names worth knowing. Group 1 is the alkali metals, soft metals that react hard with water. Group 2 is the alkaline earth metals. Group 17 is the halogens, which grab electrons eagerly. Group 18 is the noble gases, which mostly refuse to react with anything.

Groups 1, 2 and 13 through 18 are the main group elements. Groups 3 through 12 hold the transition metals, the familiar working metals like iron, copper and nickel. Two long rows are pulled out and printed underneath, which keeps the page from stretching too wide. They are the lanthanides, elements 57 through 71, and the actinides, elements 89 through 103. Slide them back into period 6 and period 7 and the table becomes one sheet 32 columns across.

Metals, nonmetals and the stair step

A jagged step runs down the right side of the table. Metals sit to the left of it. They shine, they bend rather than shatter, and they carry heat and current well. Nonmetals sit to the right. They are dull, brittle when solid, and poor conductors. Sitting right on the step are the metalloids, which act partly like each side. Boron, silicon, germanium, arsenic, antimony and tellurium are the usual six. Silicon's halfway conducting is exactly what computer chips need. Books draw the step in slightly different places, so polonium and astatine get sorted either way.

Hydrogen is printed above group 1, but it is a nonmetal and does its own thing. Of the 118 known elements, the first 94 turn up in nature, a few only in traces. The rest are built one atom at a time in accelerators and vanish in moments. At room temperature most elements are solid. Eleven are gases. Only two are liquid, mercury and bromine.

The whole table. Metals fill most of it. Nonmetals crowd the top right. The six metalloids trace the step between them. The two dashed cells mark where the lanthanide and actinide rows really belong.

Worked example · Placing an element from its atomic number

Where does selenium, Z = 34, sit? Is it a metal?

Step 1 · Count off the periodsPeriod 1 holds 2 elements, so it ends at 2. Periods 2 and 3 hold 8 each, ending at 10 and 18. Period 4 holds 18, ending at 36. Since 18 < 34 ≤ 36, selenium is in period 4.

Step 2 · Count across that periodPeriod 4 starts at Z = 19 in group 1 and fills all 18 groups in order. So the group is 34 − 18 = group 16.

Step 3 · Read off the classGroup 16, period 4 sits to the right of the step, so selenium is a nonmetal.

CheckGroup 16 also holds oxygen and sulfur, both nonmetals. Selenium belongs with them, and it does act like a heavier sulfur.

Key points

Periods are the 7 rows. Groups are the 18 columns.
One group shares outer electron count, which is why it shares behavior.
Group 1 alkali metals, group 2 alkaline earth metals, group 17 halogens, group 18 noble gases.
Groups 3 to 12 are transition metals. The lanthanides and actinides are the two rows printed below.
Metals sit left of the step, nonmetals right of it, and six metalloids sit on it.

Interactive Explore groups, periods, metals, nonmetals and metalloids. Open full ↑

Section 4Ions and ionic compounds

Move an electron and you make a charge

An atom with equal protons and electrons has no net charge. Pull an electron away and the protons now outvote the electrons, so the particle is positive. Push an extra electron on and the particle is negative. A charged atom, or a charged group of atoms, is an ion. A positive ion is a cation. A negative ion is an anion. Losing one electron gives 1+. Losing two gives 2+. Gaining one gives 1−. Note that protons never move in ordinary chemistry. Only electrons do.

The group number predicts the charge for most main group elements. Group 1 forms 1+ ions. Group 2 forms 2+. Aluminum, in group 13, forms 3+. Group 15 often takes on three electrons for 3−. Group 16 takes two for 2−. Group 17 takes one for 1−. Noble gases usually form no ion at all. Every one of those ions ends up with the same outer electron count as a noble gas, which is the pattern behind the rule.

Transition metals are less tidy. Many of them form more than one ion. Iron forms Fe²⁺ and Fe³⁺. Copper forms Cu⁺ and Cu²⁺. A handful are fixed and easy. Zinc is always Zn²⁺ and silver is always Ag⁺.

Ions built from several atoms

A polyatomic ion is a bonded group of atoms carrying one charge for the whole group. Sulfate is SO₄^2-. Its sulfur and four oxygens travel together as a single unit, and the 2− belongs to the unit, not to any one atom. Ammonium, NH₄⁺, is the common positive one. Nitrate is NO₃^- and nitrite is NO₂^-. In pairs like that, the -ate form always holds one more oxygen than the -ite form.

Ion	Formula	Ion	Formula
ammonium	NH₄⁺	hydroxide	OH^-
nitrate	NO₃^-	nitrite	NO₂^-
sulfate	SO₄^2-	sulfite	SO₃^2-
carbonate	CO₃^2-	hydrogen carbonate	HCO₃^-
phosphate	PO₄^3-	acetate	C₂H₃O₂^-
chlorate	ClO₃^-	perchlorate	ClO₄^-
cyanide	CN^-	permanganate	MnO₄^-
chromate	CrO₄^2-	dichromate	Cr₂O₇^2-

Balancing the charge writes the formula

An ionic compound pairs cations with anions, and it is usually a metal with a nonmetal. There is no single molecule to point at. Instead the ions stack into a repeating three-dimensional grid called a lattice, which runs on for as long as the crystal does. So the formula gives the smallest whole-number ratio that cancels the charge. That ratio is called a formula unit.

The rule is one line. Total positive charge plus total negative charge must come to zero. Find the smallest counts that pull that off. One useful trick is the lowest common multiple of the two charge sizes.

Left: two 3+ ions and three 2− ions both reach 6, so the charges cancel and the formula unit is Al₂O₃. Right: in a lattice every ion is ringed by ions of the opposite charge, and the grid repeats in all directions.

Worked example · Two ionic formulas from charge alone

Write the formula for aluminum oxide, and for calcium phosphate.

Step 1 · Get the two chargesAluminum is group 13, so Al³⁺. Oxygen is group 16, so O^2-.

Step 2 · Find the lowest common multipleThe charge sizes are 3 and 2, and their lowest common multiple is 6. So you need 6+ of positive and 6− of negative.

Step 3 · Divide to get the counts6 ÷ 3 = 2 aluminum ions, and 6 ÷ 2 = 3 oxide ions. The formula is Al₂O₃.

Step 4 · Repeat with a polyatomic ionCalcium is Ca²⁺ and phosphate is PO₄^3-. The charge sizes 2 and 3 give 6 again. That needs 3 calcium ions and 2 phosphate ions, so the formula is Ca₃(PO₄)₂. The parentheses keep the 2 from touching the oxygen count.

Check3(2+) = 6+ and 2(3−) = 6−, which sum to zero. Also check that the counts share no common factor, because the formula must be the smallest ratio.

Key points

Only electrons move. Losing them makes a cation, gaining them makes an anion.
Main group charges follow the group: 1+, 2+, 3+, then 3−, 2−, 1−.
Many transition metals form more than one charge, so the charge has to be stated.
A polyatomic ion carries one charge for the whole bonded group.
An ionic formula is the smallest whole-number ratio that makes the total charge zero.

Interactive Balance charges to build neutral ionic formulas. Open full ↑

Section 5Molecular compounds and formulas

Sharing instead of handing over

Two nonmetals do not pass electrons back and forth. Neither one gives up easily, so they share instead. A shared pair of electrons is a covalent bond. A neutral group of atoms held together by covalent bonds is a molecule. A molecular compound is built out of molecules, so it comes in countable units. One water molecule is two hydrogens and one oxygen, and the next one is the same.

That is a real difference from an ionic compound. Table salt has no smallest salt particle, only a ratio that repeats through the crystal. Sugar does have one. You can point at a single sucrose molecule and say that is the whole thing. Molecular compounds also tend to melt at lower temperatures and dissolve without splitting into ions.

Some elements travel as molecules of themselves. Seven are diatomic under normal conditions: H₂, N₂, O₂, F₂, Cl₂, Br₂ and I₂. Phosphorus goes around as P₄ and sulfur as S₈. A lone oxygen atom is rare and very reactive, which is why the air holds O₂ and not O.

Three ways to write the same compound

A molecular formula counts every atom in one molecule. Glucose is C₆H₁₂O₆, so six carbons, twelve hydrogens and six oxygens. An empirical formula gives the smallest whole-number ratio of those atoms. Divide every subscript in glucose by 6 and you get CH₂O. Something is lost on the way. Formaldehyde really is CH₂O, and acetic acid is C₂H₄O₂. All three share one ratio, and all three are completely different substances.

A structural formula shows which atom is bonded to which. Lines stand in for bonds. That is often the detail that decides behavior. Ethanol and dimethyl ether are both C₂H₆O. Ethanol is the alcohol in drinks. Dimethyl ether is a gas used as a propellant. Compounds that share a molecular formula and differ in structure are isomers.

Models push further still. A ball-and-stick model keeps the bonds visible so you can see the angles between them. A space-filling model drops the sticks and swells the balls, so you see the room the electrons actually occupy. Each view hides something, which is why chemists keep all of them.

One compound, three views. The molecular formula counts atoms. The empirical formula gives the ratio and loses the count. The structural formula is the only one that tells you the double bond sits on the second carbon.

Worked example · Molecular formula to empirical formula

Give the empirical formula for octane, C₈H₁₈, and for N₂O₄.

Step 1 · List the subscriptsOctane gives 8 and 18.

Step 2 · Find the largest number that divides bothFactors of 8 are 1, 2, 4 and 8. Factors of 18 are 1, 2, 3, 6, 9 and 18. The largest shared factor is 2.

Step 3 · Divide and write it8 ÷ 2 = 4 and 18 ÷ 2 = 9, so the empirical formula is C₄H₉.

Step 4 · Do the second oneN₂O₄ gives 2 and 4, which share a factor of 2. Dividing gives NO₂.

CheckIn each answer the subscripts share no factor above 1, so each is truly the smallest ratio. Note that NO₂ is also a real compound in its own right, which is exactly why an empirical formula alone cannot name a substance.

Key points

Nonmetals share electrons in covalent bonds and build molecules.
A molecular compound has a smallest unit. An ionic compound has only a repeating ratio.
Seven elements are diatomic, plus P₄ and S₈.
Molecular formula counts atoms. Empirical formula gives the smallest ratio. Structural formula shows the bonds.
Isomers share a molecular formula and differ in structure, so they differ in behavior.

Interactive Read empirical, molecular and structural formulas of the same compound. Open full ↑

Section 6Naming compounds

Two questions sort every formula

Chemists needed one naming system so that a formula reads the same in any country. The system splits by compound type, so start by deciding which type you have. Ask two questions in order. Does the formula start with hydrogen and act as an acid in water? Then name it as an acid. If not, does it start with a metal or with ammonium? If yes it is ionic, and if no it is molecular. Two questions, three roads.

Pick the road first. The same two questions sort any formula into one of three naming rules. Getting the type right is most of the work, because each road has its own small set of rules.

Ionic names

Name the cation, then the anion, and never use prefixes. The metal keeps its ordinary name. A single-atom anion drops its ending and picks up -ide. So NaCl is sodium chloride and MgO is magnesium oxide. A polyatomic ion keeps whatever name it already has, so Na₂SO₄ is sodium sulfate.

If the metal can form more than one charge, you have to say which one. Put the charge in Roman numerals in parentheses right after the metal name. FeCl₂ is iron(II) chloride and FeCl₃ is iron(III) chloride. You read that charge off the formula by balancing, never by guessing. Metals with only one common charge, such as sodium or zinc, get no numeral at all. One more case: a hydrate traps water inside its crystal, and a prefix counts it. CuSO₄·5H₂O is copper(II) sulfate pentahydrate.

Molecular names

Here you name the subscripts out loud with prefixes. Mono is 1, di is 2, tri is 3, tetra is 4, penta is 5, hexa is 6, hepta is 7, octa is 8, nona is 9 and deca is 10. The first element keeps its name and skips mono. The second element takes the -ide ending and always gets its prefix. CO is carbon monoxide and CO₂ is carbon dioxide. SF₆ is sulfur hexafluoride. Drop the final vowel of a prefix before oxide, which is why you say monoxide and pentoxide.

Acid names

Acids come in two families. A binary acid is hydrogen plus one other element. Use hydro-, then the root of that element, then -ic acid. HCl in water is hydrochloric acid. An oxyacid is hydrogen plus a polyatomic ion that holds oxygen. Here the ending of the anion steers the name. An -ate anion gives -ic acid. An -ite anion gives -ous acid. So HNO₃ is nitric acid and HNO₂ is nitrous acid. H₂SO₄ is sulfuric acid and H₂SO₃ is sulfurous acid.

Anion	Anion formula	Acid	Acid name
chloride	Cl^-	HCl	hydrochloric acid
sulfide	S^2-	H₂S	hydrosulfuric acid
nitrate	NO₃^-	HNO₃	nitric acid
nitrite	NO₂^-	HNO₂	nitrous acid
sulfate	SO₄^2-	H₂SO₄	sulfuric acid
sulfite	SO₃^2-	H₂SO₃	sulfurous acid
phosphate	PO₄^3-	H₃PO₄	phosphoric acid
perchlorate	ClO₄^-	HClO₄	perchloric acid

Why sulfuric and not sulfic? A few roots keep an extra syllable so the name stays easy to say. Sulfate gives sulfuric, and phosphate gives phosphoric. The rule still holds underneath. An -ate ion gives -ic acid, and an -ite ion gives -ous acid.

Worked example · Naming three compounds

Name Fe₂O₃, N₂O₅ and H₂SO₃.

Step 1 · Sort each oneIron is a metal, so the first is ionic. Nitrogen and oxygen are both nonmetals, so the second is molecular. The third starts with hydrogen and is an acid.

Step 2 · Work out the iron chargeOxide is always O^2-, and three of them give 3 × (2−) = 6−. The two iron ions must supply 6+, so each one is 6 ÷ 2 = 3+. That makes it iron(III) oxide.

Step 3 · Count the molecular oneTwo nitrogens give di, and five oxygens give penta. Drop the final a before oxide. That makes it dinitrogen pentoxide.

Step 4 · Follow the acid endingThe anion here is sulfite, SO₃^2-. An -ite ion gives an -ous acid, so this is sulfurous acid.

CheckRun each name backwards. Iron(III) oxide pairs Fe³⁺ with O^2-, which balances at Fe₂O₃. Dinitrogen pentoxide rebuilds N₂O₅. Sulfurous acid needs two hydrogens for a 2− ion, giving H₂SO₃.

Key points

Sort the formula first: acid, ionic or molecular. Each type has its own rules.
Ionic names use no prefixes. The anion ends in -ide unless it is polyatomic.
A metal with more than one charge needs a Roman numeral, read off by balancing.
Molecular names use prefixes on both elements, except mono on the first.
Binary acids are hydro...-ic. For oxyacids, -ate gives -ic acid and -ite gives -ous acid.

Interactive Name ionic, molecular and acid compounds step by step. Open full ↑

The chapter in one breath

Three particles build everything. Protons and neutrons pack the nucleus, and electrons move outside it.
Protons are the name tag. The atomic number Z counts them, and the mass number A adds the neutrons on.
Isotopes differ only in neutrons. They act the same in reactions, because they hold the same electrons.
The table mass is a weighted average. Multiply each isotope mass by its share, then add.
The table is a behavior map. Columns are groups and share behavior. Rows are periods.
Ions are atoms that gained or lost electrons. Group number predicts the charge for most main group elements.
An ionic formula is a charge balance. Find the smallest counts that sum to zero charge.
Molecules share electrons instead. Molecular, empirical and structural formulas each tell you a different thing.
Naming has three roads. Decide acid, ionic or molecular first, then apply that road's rules.

ReferenceWords worth knowing

Every term below showed up in bold somewhere above. Come back here whenever a word stops making sense.

alkali metals: Group 1 metals. Soft, very reactive, and they form 1+ ions.
anion: An ion with a negative charge, made by gaining one or more electrons.
atomic mass unit (u): A mass unit equal to 1/12 the mass of a carbon-12 atom, or 1.6605 × 10^-24 g.
atomic number (Z): The number of protons in a nucleus. It decides which element the atom is.
average atomic mass: The isotope masses of an element weighted by their abundances and added. This is the value printed on the table.
cation: An ion with a positive charge, made by losing one or more electrons.
covalent bond: A pair of electrons shared between two atoms.
empirical formula: The smallest whole-number ratio of atoms in a compound, such as CH₂O.
formula unit: The smallest whole-number ratio of ions that makes an ionic compound neutral.
group: A column of the periodic table, numbered 1 to 18. Members behave alike.
ion: An atom or bonded group of atoms that carries a net electric charge.
ionic compound: A compound of cations and anions held in a lattice, usually a metal with a nonmetal.
isotopes: Atoms of one element with different neutron counts, so the same Z and different A.
mass number (A): Protons plus neutrons in one nucleus. Always a whole number.
metalloid: An element on the step between metals and nonmetals that behaves partly like each.
molecular formula: The actual count of each kind of atom in one molecule, such as C₆H₁₂O₆.
oxyacid: An acid of hydrogen plus a polyatomic ion containing oxygen, such as HNO₃.
percent abundance: The share of atoms in a natural sample that are one particular isotope.
period: A row of the periodic table, numbered 1 to 7.
polyatomic ion: A bonded group of atoms carrying one charge for the whole group, such as SO₄^2-.

Check yourself

1. An atom holds 15 protons and 16 neutrons. Name it and give its mass number.

Phosphorus-31. Z = 15 names it phosphorus. A = 15 + 16 = 31. The neutron count never changes the name.

2. A particle holds 20 protons, 20 neutrons and 18 electrons. Write its symbol with charge.

⁴⁰Ca²⁺ Z = 20 is calcium. A = 20 + 20 = 40. Charge = 20 − 18 = 2+, so two electrons are missing.

3. Copper is 69.15% copper-63 at 62.9296 u and 30.85% copper-65 at 64.9278 u. Find its average atomic mass.

63.55 u. 0.6915 × 62.9296 u = 43.5158 u. 0.3085 × 64.9278 u = 20.0302 u. Adding gives 63.5460 u, so 63.55 u. It leans toward copper-63, which is the common one.

4. Boron is 19.9% boron-10 at 10.0129 u and 80.1% boron-11 at 11.0093 u. Find the average atomic mass.

10.81 u. 0.199 × 10.0129 u = 1.9926 u. 0.801 × 11.0093 u = 8.8184 u. The sum is 10.8110 u, which rounds to 10.81 u. Here the heavier isotope is the common one, so the answer sits near 11.

5. Which is larger for carbon-13, the mass number or the atomic mass? Why are they not the same thing?

The atomic mass, 13.00335 u, is a little larger than the mass number of 13. A mass number is a count of particles, so it is always a whole number. An atomic mass is a measured mass in u. Only carbon-12 has the two numbers equal, because the unit was defined from it.

6. Give the group and period of selenium, Z = 34, and say whether it is a metal.

Group 16, period 4, and it is a nonmetal. Period 4 runs from Z = 19 to Z = 36, so 34 lands there. Counting across gives 34 − 18 = group 16, the same group as oxygen and sulfur.

7. Sort these as metal, nonmetal or metalloid: Si, Ca, Ne, Sb.

Si metalloid, Ca metal, Ne nonmetal, Sb metalloid. Silicon and antimony both sit on the step. Calcium is a group 2 metal. Neon is a noble gas, far to the right of the step.

8. Write the formula for the compound of Mg²⁺ and N^3-.

Mg₃N₂ The charge sizes are 2 and 3, and their lowest common multiple is 6. You need 6 ÷ 2 = 3 magnesium ions and 6 ÷ 3 = 2 nitride ions. Check: 3(2+) = 6+ and 2(3−) = 6−, which cancel.

9. Write the formula for iron(III) sulfate.

Fe₂(SO₄)₃ The numeral says Fe³⁺, and sulfate is SO₄^2-. Sizes 3 and 2 give 6, so you need 2 iron ions and 3 sulfate ions. Check: 2(3+) = 6+ and 3(2−) = 6−. The parentheses keep the 3 from changing the oxygen count.

10. Give the empirical formula of C₆H₆ and of C₄H₁₀.

CH and C₂H₅. For C₆H₆ the largest shared factor of 6 and 6 is 6, and 6 ÷ 6 = 1 for each, giving CH. For C₄H₁₀ the largest shared factor of 4 and 10 is 2, and 4 ÷ 2 = 2 with 10 ÷ 2 = 5, giving C₂H₅.

11. Name these: (a) K₂O, (b) Cu(NO₃)₂, (c) SF₆, (d) HBr in water.

(a) potassium oxide, (b) copper(II) nitrate, (c) sulfur hexafluoride, (d) hydrobromic acid. In (b) the two nitrate ions carry 2− in total, so copper must be 2+. In (c) both parts are nonmetals, so prefixes apply. In (d) a binary acid takes hydro... plus -ic acid.

12. Write formulas for: (a) ammonium carbonate, (b) dinitrogen pentoxide, (c) phosphoric acid.

(a) (NH₄)₂CO₃, (b) N₂O₅, (c) H₃PO₄. In (a) two 1+ ammonium ions balance one 2− carbonate. In (b) the prefixes give the subscripts directly. In (c) phosphoric comes from phosphate, PO₄^3-, so three hydrogens are needed.

Previous← Chapter 1 NextChapter 3 →