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Chemistry prerequisites · Study sheet

How to balance chemical equations, molar mass, bonds and solubility rules

The chemistry a prerequisite course keeps asking for, in one place: set coefficients and watch the atom count balance, then molar mass, types of chemical reactions, ionic vs covalent bonds, a solubility rules chart, significant figures and a printable periodic table with names, with 34 flashcards and 20 practice questions. Free, and for study only.

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1 · Balance an equation

Set the coefficients and watch the atoms

Pick an equation, then use the plus and minus buttons to set the number in front of each formula. The table counts the atoms of each element on each side. When every row matches, the equation is balanced. The steps and worked examples are just below.

Choose an equation

H₂ + O₂ → H₂O

  • Reactant H₂1
  • Reactant O₂1
  • Product H₂O1
Atoms on each side
ElementLeftRightMatch
H22Equal
O21Not equal

Not balanced yet: O is 2 on the left and 1 on the right.

2 · How to balance chemical equations

The steps, with worked examples

A balanced equation has the same number of atoms of each element on both sides. Here is the routine, then four worked examples with the atom count at every step.

What an equation says

Reactants and products
Reactants are the substances that react and sit on the left. Products are the substances formed and sit on the right.
Coefficient
The number in front of a formula. It gives the relative number of that species, and a coefficient of 1 is normally left out.
Why balance at all
A balanced equation has the same number of atoms of each element on both sides, which agrees with the law of conservation of matter.
Coefficients are ratios
CH₄ + 2O₂ → CO₂ + 2H₂O says 1 : 2 : 1 : 2. It reads the same for single molecules, for dozens or for moles.
State symbols
(s) solid, (l) liquid, (g) gas and (aq) dissolved in water. A Greek capital delta over the arrow can mark a reaction carried out by heating.

The six steps

1. Write the unbalanced equation
Put the formulas of the reactants on the left and the products on the right, with plus signs between formulas and an arrow between the sides.
2. Count the atoms of each element on each side
For one formula, multiply the coefficient by the subscript. If an element is in more than one formula on a side, add the counts together.
3. Change coefficients only, never subscripts
A subscript is part of the substance’s identity. Changing H₂O to H₂O₂ would balance the atoms but give hydrogen peroxide instead of water.
4. Work one element at a time, and leave the awkward one for last
Start with an element that is in only one formula on each side. Oxygen, which often appears in several formulas, is usually best left until the end.
5. Recount after every change
Changing a coefficient upsets the counts of every element in that formula, so check all the elements again.
6. Clear fractions, then reduce
A fraction is allowed as a stepping stone. Multiply every coefficient by the same whole number to clear it, then divide by any common factor to get the smallest whole numbers.

Worked examples

Water splits into hydrogen and oxygen

  1. Step 1 · Write it unbalanced

    H₂O → H₂ + O₂

    Atoms on each side for H₂O → H₂ + O₂
    ElementLeftRightMatch
    H22Equal
    O12Not equal

    Hydrogen matches (2 and 2) but oxygen does not (1 on the left, 2 on the right).

  2. Step 2 · Fix oxygen with a coefficient

    2H₂O → H₂ + O₂

    Atoms on each side for 2H₂O → H₂ + O₂
    ElementLeftRightMatch
    H42Not equal
    O22Equal

    A 2 in front of H₂O gives 2 oxygen on each side, but now hydrogen is 4 on the left and 2 on the right.

  3. Step 3 · Fix hydrogen

    2H₂O → 2H₂ + O₂

    Atoms on each side for 2H₂O → 2H₂ + O₂
    ElementLeftRightMatch
    H44Equal
    O22Equal

    A 2 in front of H₂ makes 4 hydrogen on each side. Both elements match: balanced.

Nitrogen and oxygen make dinitrogen pentoxide

  1. Step 1 · Write it unbalanced

    N₂ + O₂ → N₂O₅

    Atoms on each side for N₂ + O₂ → N₂O₅
    ElementLeftRightMatch
    N22Equal
    O25Not equal

    Nitrogen matches. Oxygen is 2 on the left and 5 on the right.

  2. Step 2 · Aim for 10 oxygen on each side

    N₂ + 5O₂ → 2N₂O₅

    Atoms on each side for N₂ + 5O₂ → 2N₂O₅
    ElementLeftRightMatch
    N24Not equal
    O1010Equal

    10 is the lowest common multiple of 2 and 5. Oxygen now matches, but nitrogen is 2 on the left and 4 on the right.

  3. Step 3 · Fix nitrogen

    2N₂ + 5O₂ → 2N₂O₅

    Atoms on each side for 2N₂ + 5O₂ → 2N₂O₅
    ElementLeftRightMatch
    N44Equal
    O1010Equal

    A 2 in front of N₂ makes 4 nitrogen on each side: balanced.

Ethane burns: an odd number of oxygen atoms

  1. Step 1 · Balance carbon and hydrogen first

    C₂H₆ + O₂ → 3H₂O + 2CO₂

    Atoms on each side for C₂H₆ + O₂ → 3H₂O + 2CO₂
    ElementLeftRightMatch
    C22Equal
    H66Equal
    O27Not equal

    3 H₂O gives 6 hydrogen and 2 CO₂ gives 2 carbon. That leaves 7 oxygen atoms on the right, an odd number.

  2. Step 2 · Use a fraction for now

    C₂H₆ + 7⁄2O₂ → 3H₂O + 2CO₂

    Atoms on each side for C₂H₆ + 7⁄2O₂ → 3H₂O + 2CO₂
    ElementLeftRightMatch
    C22Equal
    H66Equal
    O77Equal

    No whole number of O₂ gives 7 oxygen atoms, so use 7⁄2 as a temporary coefficient.

  3. Step 3 · Multiply every coefficient by 2

    2C₂H₆ + 7O₂ → 6H₂O + 4CO₂

    Atoms on each side for 2C₂H₆ + 7O₂ → 6H₂O + 4CO₂
    ElementLeftRightMatch
    C44Equal
    H1212Equal
    O1414Equal

    Doubling keeps the balance and clears the fraction: 2C₂H₆ + 7O₂ → 6H₂O + 4CO₂.

Balanced, but not in the smallest numbers

  1. Step 1 · This is balanced

    3N₂ + 9H₂ → 6NH₃

    Atoms on each side for 3N₂ + 9H₂ → 6NH₃
    ElementLeftRightMatch
    N66Equal
    H1818Equal

    Every element matches, but the coefficients share a common factor of 3.

  2. Step 2 · Divide by the common factor

    N₂ + 3H₂ → 2NH₃

    Atoms on each side for N₂ + 3H₂ → 2NH₃
    ElementLeftRightMatch
    N22Equal
    H66Equal

    The convention is the smallest whole-number coefficients: N₂ + 3H₂ → 2NH₃.

Molecular, complete ionic and net ionic equations

When ionic compounds react in water, the same reaction can be written at three levels of detail. This is the textbook’s example, calcium chloride mixed with silver nitrate.

  1. Molecular equation

    CaCl₂(aq) + 2AgNO₃(aq) → Ca(NO₃)₂(aq) + 2AgCl(s)

    Writes each compound as one formula, without showing the ions.

  2. Complete ionic equation

    Ca²⁺(aq) + 2Cl⁻(aq) + 2Ag⁺(aq) + 2NO₃⁻(aq) → Ca²⁺(aq) + 2NO₃⁻(aq) + 2AgCl(s)

    Shows every dissolved ionic compound as its separate ions. The solid AgCl stays together.

  3. Net ionic equation

    Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

    Drops the spectator ions, Ca²⁺ and NO₃⁻, which are the same on both sides, and uses the smallest whole numbers.

3 · Molar mass and the mole

How to find molar mass

Add up the average atomic masses of every atom in the formula. The total in grams per mole is the molar mass.

  1. Write the formula and list each element with how many atoms of it there are (for brackets, multiply what is inside by the number outside).

  2. Multiply each element’s atom count by its average atomic mass.

  3. Add the results. The total is the formula mass in amu and the molar mass in g/mol.

  4. To go from grams to moles, divide the mass by the molar mass. To go from moles to grams, multiply.

Vocabulary

Formula mass
The sum of the average atomic masses of all the atoms in a formula, in amu. For a molecular compound it can be called the molecular mass.
Mole
An amount unit, like a dozen, for counting atoms or molecules. One mole is 6.022 × 10²³ particles (Avogadro’s number, here rounded).
Molar mass
The mass in grams of one mole of a substance, in g/mol. It is numerically equal to the formula mass in amu.
Same count, different mass
One mole of any element has the same number of atoms as one mole of any other, but the masses differ because individual atoms differ in mass.

Worked molar masses

Atomic masses used: H 1.008, C 12.01, N 14.01, O 16.00, Na 22.99, P 30.97, Cl 35.45, Ca 40.08 g/mol.

  • Water

    H₂O

    2 × 1.008 + 16.00. 2 H and 1 O.

    18.02 g/mol

  • Carbon dioxide

    CO₂

    12.01 + 2 × 16.00. 1 C and 2 O.

    44.01 g/mol

  • Sodium chloride (table salt)

    NaCl

    22.99 + 35.45. An ionic compound, so this is a formula mass, not a molecular mass.

    58.44 g/mol

  • Glycine

    C₂H₅O₂N

    2 × 12.01 + 5 × 1.008 + 2 × 16.00 + 14.01. 2 C, 5 H, 2 O and 1 N.

    75.07 g/mol

  • Glucose

    C₆H₁₂O₆

    6 × 12.01 + 12 × 1.008 + 6 × 16.00. 6 C, 12 H and 6 O.

    180.16 g/mol

  • Calcium phosphate

    Ca₃(PO₄)₂

    3 × 40.08 + 2 × 30.97 + 8 × 16.00. The 2 outside the brackets doubles P and O: 3 Ca, 2 P and 8 O.

    310.18 g/mol

Going between grams and moles: moles = grams ÷ molar mass. For 28.35 g of glycine at 75.07 g/mol, that is 28.35 ÷ 75.07 = 0.3776 mol. Going the other way, grams = moles × molar mass.

4 · Types of chemical reactions

The types of chemical reactions, with an example of each

Courses name reactions by pattern. Chemistry 2e groups the common ones as precipitation, acid-base and oxidation-reduction (redox) reactions, with combustion and single replacement as kinds of redox.

  • Synthesis (combination)

    Pattern: A + B → AB

    2Na(s) + Cl₂(g) → 2NaCl(s)

    Two or more reactants join to make one product. Chemistry 2e also classes this one as a redox reaction.

  • Decomposition

    Pattern: AB → A + B

    2H₂O(l) → 2H₂(g) + O₂(g)

    One reactant breaks into simpler substances. Heating is often written as Δ over the arrow, as in CaCO₃(s) → CaO(s) + CO₂(g).

  • Single replacement (single displacement)

    Pattern: A + BC → AC + B

    Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)

    An element takes the place of an ion in a compound. In Chemistry 2e these are redox reactions in which a metal is oxidized.

  • Double replacement (precipitation)

    Pattern: AB + CD → AD + CB

    2KI(aq) + Pb(NO₃)₂(aq) → PbI₂(s) + 2KNO₃(aq)

    Dissolved ionic compounds swap ions. In a precipitation reaction one product is a solid; also called double displacement or metathesis.

  • Acid-base (neutralization)

    Pattern: acid + base → salt + water

    Mg(OH)₂(s) + 2HCl(aq) → MgCl₂(aq) + 2H₂O(l)

    An acid-base reaction is one in which H⁺ is transferred. Neutralization is the case where the reactants are an acid and a base.

  • Combustion

    Pattern: fuel + oxidant → products, with heat

    CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l)

    A fuel (the reductant) reacts vigorously with an oxidant, often oxygen, giving off heat and often light as a flame. It is a kind of redox reaction.

Oxidation and reduction

Oxidation
Loss of electrons, which is an increase in oxidation number.
Reduction
Gain of electrons, which is a decrease in oxidation number.
Reducing agent (reductant)
The species that is oxidized. It supplies the electrons.
Oxidizing agent (oxidant)
The species that is reduced. It takes the electrons.
Example
In 2Na + Cl₂ → 2NaCl, sodium goes from 0 to +1 (oxidized, the reducing agent) and chlorine from 0 to −1 (reduced, the oxidizing agent).

Acids and bases

Acid
In this chapter’s sense, a substance that dissolves in water to give hydronium ions, H₃O⁺.
Base
A substance that dissolves in water to give hydroxide ions, OH⁻. The most common are group 1 or group 2 metal hydroxides such as NaOH and Ca(OH)₂.
Strong acids in the textbook list
HBr, HCl, HI, HNO₃, HClO₄ and H₂SO₄ react completely with water.

5 · Ionic vs covalent bonds

Ionic vs covalent bonds

Two ways atoms bond: ions attracting each other, or atoms sharing electrons. Compared side by side, then the vocabulary for polar and nonpolar bonds.

  • What holds it together

    Ionic: The electrostatic attraction between oppositely charged ions, a cation (+) and an anion (−).

    Covalent: Two atoms attracting a shared pair of electrons.

  • Which atoms

    Ionic: Typically a metal (forms the cation) with a nonmetal (forms the anion).

    Covalent: Frequently a nonmetal with another nonmetal.

  • What the electrons do

    Ionic: Typically transferred from one atom to the other.

    Covalent: Shared between the two atoms.

  • The unit

    Ionic: A lattice of many ions. The formula gives the simplest ratio, so NaCl is a formula unit, not a “molecule”.

    Covalent: A molecule. The formula shows the atoms in one molecule.

  • Melting and boiling points

    Ionic: Tend to be high. Ionic solids are rigid and brittle.

    Covalent: Generally much lower. Many are liquids or gases at room temperature.

  • Conducting electricity

    Ionic: Poor as a solid. Excellent once dissolved or melted, because the ions can move.

    Covalent: Most are poor conductors in any state.

  • In water

    Ionic: Most ionic solids dissolve readily.

    Covalent: Most covalent compounds are insoluble in water.

  • Examples

    Ionic: NaCl, Al₂O₃ (2 Al³⁺ and 3 O²⁻).

    Covalent: H₂, Cl₂, HCl.

Polar, nonpolar and electronegativity

Pure (nonpolar) covalent bond
A bond between identical atoms, as in H₂ and Cl₂. The electrons are shared equally.
Polar covalent bond
A bond between different atoms. The electrons are shared unequally, so one atom has a partial negative charge (δ−) and the other a partial positive charge (δ+).
Electronegativity
A measure of how strongly an atom attracts the electrons in a bond. It generally rises from left to right across a period and falls down a group; fluorine is the highest, at 4.0 on Pauling’s scale.
Electronegativity difference as a guide
The bigger the difference between two bonded atoms, the more the bond looks ionic. The textbook gives H–H as 0 (nonpolar), H–Cl as 0.9 (polar covalent) and Na–Cl as 2.1 (ionic), and calls this a rough guide with exceptions.
Both kinds in one compound
In potassium nitrate, KNO₃, the K⁺ and NO₃⁻ ions are held by an ionic bond, while the atoms inside NO₃⁻ are held by polar covalent bonds.

Ion charges and writing ionic formulas

Groups 1 and 2
The group number is the charge of the cation: Na⁺ and Ca²⁺.
Groups 13 to 17 (cations)
The charge is the group number minus 10, so aluminum in group 13 forms Al³⁺.
Nonmetal anions
The charge is the number of electrons needed to fill the outer shell. Oxygen gains two, so the oxide ion is O²⁻.
Writing the formula
The total positive charge must equal the total negative charge. Al₂O₃ has (2 × +3) + (3 × −2) = 0.

6 · Solubility rules

Solubility rules chart

Patterns for which ionic compounds dissolve in water, used to predict whether mixing two solutions makes a solid precipitate.

Soluble ionic compounds

  • Soluble if it contains: NH₄⁺ and the group 1 cations: Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺

    Exceptions: None

  • Soluble if it contains: Cl⁻, Br⁻, I⁻

    Exceptions: Compounds with Ag⁺, Hg₂²⁺ and Pb²⁺

  • Soluble if it contains: F⁻

    Exceptions: Compounds with group 2 cations, Pb²⁺ and Fe³⁺

  • Soluble if it contains: C₂H₃O₂⁻ (acetate), HCO₃⁻, NO₃⁻, ClO₃⁻

    Exceptions: None

  • Soluble if it contains: SO₄²⁻

    Exceptions: Compounds with Ag⁺, Ba²⁺, Ca²⁺, Hg₂²⁺, Pb²⁺ and Sr²⁺

Insoluble ionic compounds

  • Insoluble if it contains: CO₃²⁻, CrO₄²⁻, PO₄³⁻, S²⁻

    Exceptions: Compounds with group 1 cations and NH₄⁺

  • Insoluble if it contains: OH⁻

    Exceptions: Compounds with group 1 cations and Ba²⁺

These are textbook guidelines for predicting precipitates from reasonably concentrated solutions, not exact limits. Real solubility is a matter of degree, and your course’s chart may be set out differently, so use the chart your instructor gives you.

Predicting a precipitate

  1. List the ions present when the two solutions are mixed.

  2. Swap partners: pair each cation with the other solution’s anion.

  3. Look each new pairing up in the chart. An insoluble pairing is a predicted precipitate.

  4. If a precipitate forms, write the net ionic equation: the dissolved ions that make the solid, and nothing else.

Mixing silver nitrate and sodium chloride gives Ag⁺, NO₃⁻, Na⁺ and Cl⁻. The new pairings are NaNO₃ (all nitrates are soluble) and AgCl (a chloride with Ag⁺, an exception, so insoluble). A precipitate of AgCl is predicted: Ag⁺(aq) + Cl⁻(aq) → AgCl(s).

7 · Significant figures

Significant figures rules

Significant figures show how certain a measurement is. First which digits count, then how many to keep after adding, subtracting, multiplying or dividing.

Which digits count

Nonzero digits
Always significant.
Captive zeros (between nonzero digits)
Always significant, because they result from the measurement.
Leading zeros (before the first nonzero digit)
Never significant. They only show where the decimal point is.
Trailing zeros after the decimal point
Significant. Count from the first nonzero digit on the left to the end.
Trailing zeros in a whole number (1,300)
Ambiguous. Scientific notation settles it: 1.3 × 10³ has two, 1.30 × 10³ three and 1.300 × 10³ four. With no other information, assume they are not significant.
Exact numbers
Counted objects (eggs in a carton) and defined quantities are exact and have no uncertainty.

Counting examples

  • 0.00832407

    6 significant figures

    The three leading zeros do not count; 8, 3, 2, 4, 0 and 7 all do (the 0 is captive).

  • 6.72

    3 significant figures

    All three digits are nonzero.

  • 6.723

    4 significant figures

    All four digits are nonzero.

  • 0.051065

    5 significant figures

    The leading zeros do not count; 5, 1, 0, 6 and 5 do (the 0 is captive).

  • 5.00

    3 significant figures

    The two trailing zeros are after the decimal point, so they count.

  • 1.300 × 10³

    4 significant figures

    In scientific notation every digit in the front number counts.

Rounding in calculations

Adding and subtracting
Round the result to the same number of decimal places as the number with the fewest decimal places.
Multiplying and dividing
Round the result to the same number of significant figures as the number with the fewest.
Rounding
If the dropped digit is below 5, leave the kept digit alone. If it is above 5, round up. A dropped 5 on its own is a special case that books and instructors handle differently, so follow your course.

  • 1.0023 g + 4.383 g

    5.3853 g → 5.385 g

    Addition: the fewest decimal places is three (4.383), so round to the thousandths place.

  • 486 g − 421.23 g

    64.77 g → 65 g

    Subtraction: 486 has no decimal places, so round to the ones place.

  • 0.6238 cm × 6.6 cm

    4.11708 cm² → 4.1 cm²

    Multiplication: 6.6 has the fewest significant figures (two).

  • 421.23 g ÷ 486 mL

    0.866728… g/mL → 0.867 g/mL

    Division: 486 has three significant figures, the fewest.

8 · Periodic table

Printable periodic table with names

All 118 elements in order of atomic number, each box with the number, symbol and name. Shaded boxes are the transition metals and the two inner rows. Printing this page prints the table only; choose landscape in the print dialog for the largest boxes, or save it as a PDF from the same dialog.

Periodic table of the elements, 7 periods by 18 groups
Group 1Group 2Group 3Group 4Group 5Group 6Group 7Group 8Group 9Group 10Group 11Group 12Group 13Group 14Group 15Group 16Group 17Group 18
Period 11HHydrogen2HeHelium
Period 23LiLithium4BeBeryllium5BBoron6CCarbon7NNitrogen8OOxygen9FFluorine10NeNeon
Period 311NaSodium12MgMagnesium13AlAluminum14SiSilicon15PPhosphorus16SSulfur17ClChlorine18ArArgon
Period 419KPotassium20CaCalcium21ScScandium22TiTitanium23VVanadium24CrChromium25MnManganese26FeIron27CoCobalt28NiNickel29CuCopper30ZnZinc31GaGallium32GeGermanium33AsArsenic34SeSelenium35BrBromine36KrKrypton
Period 537RbRubidium38SrStrontium39YYttrium40ZrZirconium41NbNiobium42MoMolybdenum43TcTechnetium44RuRuthenium45RhRhodium46PdPalladium47AgSilver48CdCadmium49InIndium50SnTin51SbAntimony52TeTellurium53IIodine54XeXenon
Period 655CsCesium56BaBarium57–7172HfHafnium73TaTantalum74WTungsten75ReRhenium76OsOsmium77IrIridium78PtPlatinum79AuGold80HgMercury81TlThallium82PbLead83BiBismuth84PoPolonium85AtAstatine86RnRadon
Period 787FrFrancium88RaRadium89–103104RfRutherfordium105DbDubnium106SgSeaborgium107BhBohrium108HsHassium109MtMeitnerium110DsDarmstadtium111RgRoentgenium112CnCopernicium113NhNihonium114FlFlerovium115McMoscovium116LvLivermorium117TsTennessine118OgOganesson
Lanthanides 57–7157LaLanthanum58CeCerium59PrPraseodymium60NdNeodymium61PmPromethium62SmSamarium63EuEuropium64GdGadolinium65TbTerbium66DyDysprosium67HoHolmium68ErErbium69TmThulium70YbYtterbium71LuLutetium
Actinides 89–10389AcActinium90ThThorium91PaProtactinium92UUranium93NpNeptunium94PuPlutonium95AmAmericium96CmCurium97BkBerkelium98CfCalifornium99EsEinsteinium100FmFermium101MdMendelevium102NoNobelium103LrLawrencium

How the table is organized

The periodic law
The properties of the elements are periodic functions of their atomic numbers. The table lists the elements in order of atomic number and puts elements with similar properties in the same column.
Periods and groups
The seven horizontal rows are periods. The 18 vertical columns are groups, numbered 1 to 18 under the IUPAC recommendation.
Main-group, transition and inner transition
Groups 1, 2 and 13–18 hold the main-group (representative) elements. Groups 3–12 hold the transition metals. The two rows at the bottom are the inner transition metals: the lanthanides on top and the actinides below.
Group names
Group 1 (except hydrogen) is the alkali metals, group 2 the alkaline earth metals, group 16 the chalcogens, group 17 the halogens and group 18 the noble gases.
Metals, nonmetals, metalloids
Metals are shiny, malleable and conduct heat and electricity well. Nonmetals are dull and conduct poorly. Metalloids conduct moderately well and have some properties of each. Nonmetals sit in the upper right of the table.
What a box shows
Most periodic tables show each element’s atomic number, symbol, average atomic mass and sometimes its name. The table on this page shows the number, symbol and name, so it can print compactly.
See all 118 elements as a list
  1. 1HHydrogen
  2. 2HeHelium
  3. 3LiLithium
  4. 4BeBeryllium
  5. 5BBoron
  6. 6CCarbon
  7. 7NNitrogen
  8. 8OOxygen
  9. 9FFluorine
  10. 10NeNeon
  11. 11NaSodium
  12. 12MgMagnesium
  13. 13AlAluminum
  14. 14SiSilicon
  15. 15PPhosphorus
  16. 16SSulfur
  17. 17ClChlorine
  18. 18ArArgon
  19. 19KPotassium
  20. 20CaCalcium
  21. 21ScScandium
  22. 22TiTitanium
  23. 23VVanadium
  24. 24CrChromium
  25. 25MnManganese
  26. 26FeIron
  27. 27CoCobalt
  28. 28NiNickel
  29. 29CuCopper
  30. 30ZnZinc
  31. 31GaGallium
  32. 32GeGermanium
  33. 33AsArsenic
  34. 34SeSelenium
  35. 35BrBromine
  36. 36KrKrypton
  37. 37RbRubidium
  38. 38SrStrontium
  39. 39YYttrium
  40. 40ZrZirconium
  41. 41NbNiobium
  42. 42MoMolybdenum
  43. 43TcTechnetium
  44. 44RuRuthenium
  45. 45RhRhodium
  46. 46PdPalladium
  47. 47AgSilver
  48. 48CdCadmium
  49. 49InIndium
  50. 50SnTin
  51. 51SbAntimony
  52. 52TeTellurium
  53. 53IIodine
  54. 54XeXenon
  55. 55CsCesium
  56. 56BaBarium
  57. 57LaLanthanum
  58. 58CeCerium
  59. 59PrPraseodymium
  60. 60NdNeodymium
  61. 61PmPromethium
  62. 62SmSamarium
  63. 63EuEuropium
  64. 64GdGadolinium
  65. 65TbTerbium
  66. 66DyDysprosium
  67. 67HoHolmium
  68. 68ErErbium
  69. 69TmThulium
  70. 70YbYtterbium
  71. 71LuLutetium
  72. 72HfHafnium
  73. 73TaTantalum
  74. 74WTungsten
  75. 75ReRhenium
  76. 76OsOsmium
  77. 77IrIridium
  78. 78PtPlatinum
  79. 79AuGold
  80. 80HgMercury
  81. 81TlThallium
  82. 82PbLead
  83. 83BiBismuth
  84. 84PoPolonium
  85. 85AtAstatine
  86. 86RnRadon
  87. 87FrFrancium
  88. 88RaRadium
  89. 89AcActinium
  90. 90ThThorium
  91. 91PaProtactinium
  92. 92UUranium
  93. 93NpNeptunium
  94. 94PuPlutonium
  95. 95AmAmericium
  96. 96CmCurium
  97. 97BkBerkelium
  98. 98CfCalifornium
  99. 99EsEinsteinium
  100. 100FmFermium
  101. 101MdMendelevium
  102. 102NoNobelium
  103. 103LrLawrencium
  104. 104RfRutherfordium
  105. 105DbDubnium
  106. 106SgSeaborgium
  107. 107BhBohrium
  108. 108HsHassium
  109. 109MtMeitnerium
  110. 110DsDarmstadtium
  111. 111RgRoentgenium
  112. 112CnCopernicium
  113. 113NhNihonium
  114. 114FlFlerovium
  115. 115McMoscovium
  116. 116LvLivermorium
  117. 117TsTennessine
  118. 118OgOganesson

9 · Flashcards

34 chemistry flashcards

Balancing equations, molar mass, reaction types, bonds, solubility, significant figures and the periodic table. Tap a card (or press Enter) to flip it, use the arrows or the left and right keys to move, and shuffle when the order starts to feel familiar.

Card 1 of 34, Balancing. Prompt: Reactants vs products.

See all 34 cards as a list
  • Reactants vs productsReactants are on the left of the arrow; products are on the right.
  • What a coefficient tells youThe relative number of molecules (or moles) of that substance. A coefficient of 1 is left out.
  • Can you change a subscript to balance an equation?No. A subscript is part of the substance’s identity; only coefficients change.
  • Why does an equation have to balance?The same number of atoms of each element must appear on both sides, which follows the law of conservation of matter.
  • Which element should you usually leave for last?One that appears in more than one formula on a side, often oxygen.
  • You end up needing 7⁄2 as a coefficient. Now what?Multiply every coefficient by 2 to clear the fraction.
  • 3N₂ + 9H₂ → 6NH₃ is balanced. What is wrong with it?The coefficients share a factor of 3. The smallest whole numbers give N₂ + 3H₂ → 2NH₃.
  • Spectator ionsIons that are the same on both sides of a complete ionic equation. They are left out of the net ionic equation.
  • Molar massThe mass in grams of one mole of a substance, in g/mol; numerically the same as its formula mass in amu.
  • One mole is how many particles?6.022 × 10²³ (Avogadro’s number, rounded).
  • Grams to molesDivide the mass by the molar mass.
  • Moles to gramsMultiply the number of moles by the molar mass.
  • Molar mass of H₂O18.02 g/mol (2 × 1.008 + 16.00).
  • Molar mass of NaCl58.44 g/mol (22.99 + 35.45).
  • Acid + base →A salt and water (a neutralization reaction).
  • Precipitation reactionDissolved substances react to form a solid product. Also called double replacement.
  • Oxidation vs reductionOxidation is loss of electrons; reduction is gain of electrons.
  • CombustionA fuel and an oxidant (often O₂) react vigorously, giving off heat and often light.
  • Single replacement exampleZn + 2HCl → ZnCl₂ + H₂: an element takes the place of an ion in a compound.
  • Ionic bondThe electrostatic attraction between oppositely charged ions, typically a metal and a nonmetal.
  • Covalent bondAtoms sharing a pair of electrons, frequently two nonmetals.
  • Pure covalent vs polar covalentPure: identical atoms share equally (H₂). Polar: different atoms share unequally (HCl).
  • ElectronegativityHow strongly an atom attracts the electrons in a bond. Highest for fluorine.
  • Why NaCl is not a “molecule”Its ions form a lattice with no single pair bonded to each other, so the formula is just the simplest ratio.
  • Always soluble (no exceptions in the chart)Compounds of NH₄⁺ and the group 1 cations; also acetates, HCO₃⁻, nitrates and chlorates.
  • Chlorides, bromides and iodides: the exceptionsCompounds with Ag⁺, Hg₂²⁺ and Pb²⁺ are insoluble.
  • Carbonates, phosphates and sulfidesInsoluble, except with group 1 cations and NH₄⁺.
  • HydroxidesInsoluble, except with group 1 cations and Ba²⁺.
  • Leading zerosNever significant; they only locate the decimal point.
  • Add or subtractRound to the fewest decimal places.
  • Multiply or divideRound to the fewest significant figures.
  • Periods vs groupsPeriods are the 7 horizontal rows; groups are the 18 vertical columns.
  • Group 17 and group 18Group 17 is the halogens; group 18 is the noble gases.
  • How the modern table is orderedBy increasing atomic number, with similar elements in the same column.

10 · Practice questions

20 practice questions

Balancing, molar mass and moles, bonds, reaction types, solubility, significant figures and the periodic table, multiple choice, one at a time. Every answer is explained, and your score is free: no sign-up needed.

Question 1 of 20

Which coefficients balance H₂ + O₂ → H₂O (in that order)?

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Where the facts come from

Every fact on this page is reworded from one of these public textbook chapters. Atomic masses are rounded as in the textbook’s examples. Check anything important against your course materials.

Questions

Common questions

Write the unbalanced equation, count the atoms of each element on each side (coefficient times subscript, added over every formula that contains the element), then change coefficients, never subscripts, until every element matches. Leave an element that appears in several formulas, often oxygen, for last. If a fraction is needed along the way, multiply every coefficient by the same whole number, and finish by dividing out any common factor so the coefficients are the smallest whole numbers.

A subscript is part of what the substance is. Changing H₂O to H₂O₂ would balance the oxygen but name hydrogen peroxide instead of water. Only the coefficients, which say how many of each, are free to change.

List each element in the formula with its number of atoms (multiply what is in brackets by the number outside), multiply each count by the element’s average atomic mass, and add. The total, in grams per mole, is the molar mass. For water, 2 × 1.008 + 16.00 = 18.02 g/mol.

An ionic bond is the electrostatic attraction between oppositely charged ions, typically a metal and a nonmetal, as in NaCl. A covalent bond is two atoms sharing a pair of electrons, frequently two nonmetals, as in H₂ or HCl. Ionic compounds tend to have high melting and boiling points; covalent compounds generally have much lower ones.

Courses usually name synthesis, decomposition, single replacement, double replacement, combustion and acid-base neutralization. The textbook this page follows groups reactions as precipitation, acid-base and oxidation-reduction (redox), with combustion and single replacement as kinds of redox.

List the ions in the mixed solutions, pair each cation with the other solution’s anion, and look each new compound up in the chart. If one is insoluble, a precipitate is predicted. They are guidelines, and your course’s chart may differ in detail.

Every nonzero digit counts, zeros between nonzero digits count, leading zeros never count, and trailing zeros after a decimal point count. In a whole number such as 1,300 the trailing zeros are ambiguous unless scientific notation is used. For adding and subtracting, keep the fewest decimal places; for multiplying and dividing, keep the fewest significant figures.

Yes. The table on this page shows all 118 elements with number, symbol and name, and printing the page prints only the table. Landscape gives the biggest boxes, and the print dialog can also save it as a PDF.

No. It is a study aid that follows a free open textbook, so use it to review. Your own textbook, periodic table and instructor come first, especially for atomic masses, solubility charts and how your course rounds.

The trainer, the whole reference, the periodic table, the flashcards and the 20 practice questions with your score are free with no sign-up. An email unlocks the bonus: the printable chemistry sheet and 16 more questions.