The Chemistry Mistakes I See Every Semester (And How to Avoid Them)
I've been teaching chemistry for fifteen years. Every semester, I grade hundreds of exams. And every semester, I see the same mistakes.
These aren't hard concepts. They're usually careless errors, things students would catch if they knew what to look for. Here are the ten mistakes I see most often, and how to avoid them.
Confusing Atoms, Ions, and Isotopes
This one costs students points constantly.
Sodium atom (Na) and sodium ion (Na⁺) are not the same thing. The atom has 11 electrons. The ion has 10. The atom explodes in water. The ion is in your table salt.
Carbon-12 and Carbon-14 are the same element, not different elements. They're isotopes. Same number of protons (6), different number of neutrons. If the proton count matches, it's the same element.
Here's the rule: Protons define the element. Electrons define the charge. Neutrons define the isotope.
If you see a superscript charge (like ⁺ or ⁻), you're looking at an ion, not an atom. Different thing entirely.
Mistake #2: Not Balancing Chemical Equations
The Problem
Students write:
H₂ + O₂ → H₂O ✗ (WRONG)
Why It's Wrong
Law of Conservation of Mass: Atoms can't be created or destroyed.
Count atoms:
- Left side: 2 H, 2 O
- Right side: 2 H, 1 O
- Missing: 1 oxygen atom!
The Correct Way
2H₂ + O₂ → 2H₂O ✓ (CORRECT)
Count atoms:
- Left side: 4 H, 2 O
- Right side: 4 H, 2 O
- Balanced!
How to Avoid
Steps:
- Write the correct formulas (never change subscripts!)
- Count atoms on each side
- Add coefficients (big numbers in front) to balance
- Double-check each element
Common mistake within the mistake: Changing H₂O to H₂O₂ to balance. NO! This changes the substance (now it's hydrogen peroxide, not water).
Rule: Change coefficients (numbers in front), NEVER subscripts (numbers inside formulas).
Mistake #3: Mixing Up Moles and Grams
The Problem
Question: How many moles are in 36 grams of H₂O?
Student writes: 36 moles ✗
Why It's Wrong
Grams ≠ Moles
Think of it like currency:
- Grams = dollars
- Moles = euros
- You need an exchange rate (molar mass) to convert
The Correct Way
Step 1: Find molar mass of H₂O
- H: 1 g/mol × 2 = 2 g/mol
- O: 16 g/mol × 1 = 16 g/mol
- Total: 18 g/mol
Step 2: Convert grams to moles
36 g × (1 mol / 18 g) = 2 moles ✓
How to Avoid
Remember:
- Molar mass is your conversion factor
- Moles = grams ÷ molar mass
- Grams = moles × molar mass
Always write units in your calculation. If units don't cancel properly, you did it wrong!
Mistake #4: Forgetting Significant Figures
The Problem
Calculation: 12.5 × 2.1 = 26.25
Student writes: 26.25 ✗
Why It's Wrong
Your answer can't be more precise than your measurements!
Rule: Your answer can only have as many sig figs as your least precise measurement.
The Correct Way
12.5 (3 sig figs) × 2.1 (2 sig figs) = 26 (2 sig figs) ✓
How to Avoid
Sig Fig Rules:
Multiplication/Division: Use the FEWEST sig figs from any number
- 5.00 (3) × 2.1 (2) = 11 (2 sig figs)
Addition/Subtraction: Use the FEWEST decimal places
- 12.11 + 1.1 = 13.2 (one decimal place, matching 1.1)
Counting numbers: Infinite sig figs
- 3 moles = exactly 3 (not 3.0 or 3.00)
Tip: Do the full calculation, then round at the very end.
Mistake #5: Ignoring Charge Balance in Ionic Compounds
The Problem
Students write:
- NaCl₂ ✗
- MgO₂ ✗
- CaF ✗
Why It's Wrong
Ionic compounds must be electrically neutral.
Charges must balance:
- Na is +1, Cl is -1 → NaCl (not NaCl₂)
- Mg is +2, O is -2 → MgO (not MgO₂)
- Ca is +2, F is -1 → CaF₂ (not CaF)
The Correct Way
Criss-cross method:
- Write charges: Mg²⁺ and F⁻
- Criss-cross the numbers (ignore signs): Mg₁F₂
- Simplify if needed: MgF₂ ✓
Check: (2+) + 2(1-) = 0 ✓ (neutral)
How to Avoid
Always check:
- What are the charges of each ion?
- Do the charges balance to zero?
- Is the formula in lowest terms?
Memory trick: Positive and negative must cancel out. The compound can't have a net charge.
Mistake #6: Confusing Molecular and Empirical Formulas
The Problem
Question: What's the empirical formula of C₆H₁₂O₆?
Student writes: C₆H₁₂O₆ ✗
Why It's Wrong
Molecular formula: Actual number of each atom (C₆H₁₂O₆)
Empirical formula: Simplest whole-number ratio (CH₂O)
They asked for empirical, not molecular!
The Correct Way
C₆H₁₂O₆ → Divide all by 6 → CH₂O ✓
Both formulas represent glucose, but:
- Molecular (C₆H₁₂O₆): The real molecule
- Empirical (CH₂O): The simplest ratio
How to Avoid
Read the question!
- Empirical: Reduce to lowest terms
- Molecular: Actual molecule
Example:
- H₂O₂ (molecular) → HO (empirical)
- C₂H₆ (molecular) → CH₃ (empirical)
Mistake #7: Wrong Oxidation States
The Problem
Question: What's the oxidation state of Mn in KMnO₄?
Student: +4 ✗
Why It's Wrong
Oxidation states must add up to the total charge (0 for neutral compounds).
KMnO₄:
- K: +1
- Mn: ?
- O: -2 (× 4 = -8)
Calculation:
(+1) + (Mn) + (-8) = 0
Mn = +7 ✓
The Correct Way
Rules (in order):
- Free elements: 0 (O₂, Na, Fe)
- Group 1: +1 (Li, Na, K)
- Group 2: +2 (Mg, Ca)
- Fluorine: -1 (always)
- Oxygen: -2 (usually, except peroxides -1)
- Hydrogen: +1 (with nonmetals), -1 (with metals)
- Sum must equal charge (0 for neutral, charge for ions)
How to Avoid
Work systematically:
- Assign known oxidation states first
- Set up equation: sum = total charge
- Solve for unknown
Check: Do all oxidation states add up correctly?
Mistake #8: Using Wrong Units for Gas Laws
The Problem
Given: P = 2 atm, V = 5 L, T = 25°C
Student uses PV = nRT:
(2)(5) = n(0.0821)(25)
n = 4.88 moles ✗
Why It's Wrong
Temperature must be in Kelvin for gas laws!
25°C ≠ 25 K
25°C = 298 K
The Correct Way
Convert first:
T = 25°C + 273 = 298 K
Then calculate:
(2 atm)(5 L) = n(0.0821 L·atm/mol·K)(298 K)
n = 0.408 moles ✓
How to Avoid
Gas law checklist:
- Temperature: ALWAYS in Kelvin (K = °C + 273)
- Pressure: Match your R constant (usually atm)
- Volume: Match your R constant (usually L)
R constants:
- 0.0821 L·atm/(mol·K)
- 8.314 J/(mol·K)
- 62.36 L·torr/(mol·K)
Match units to your R!
Mistake #9: Forgetting to Include States of Matter
The Problem
Student writes:
NaCl + AgNO₃ → AgCl + NaNO₃ ✗
Why It's Incomplete
Missing states of matter:
- (s) = solid
- (l) = liquid
- (g) = gas
- (aq) = aqueous (dissolved in water)
States matter for:
- Precipitation reactions
- Entropy calculations
- Solubility predictions
- Understanding reaction mechanisms
The Correct Way
NaCl(aq) + AgNO₃(aq) → AgCl(s) + NaNO₃(aq) ✓
Now we see:
- AgCl precipitates out (solid)
- Other compounds stay dissolved
How to Avoid
Always indicate states when:
- Writing balanced equations
- Describing reactions
- Doing thermodynamics problems
Check solubility rules to know what precipitates!
Mistake #10: Mixing Up Endothermic and Exothermic
The Problem
Question: Is melting ice endothermic or exothermic?
Student: Exothermic because it releases cold ✗
Why It's Wrong
"Releases cold" isn't a thing in chemistry!
Energy perspective:
- Endothermic: Absorbs heat (ΔH positive)
- Exothermic: Releases heat (ΔH negative)
Melting ice:
- Ice absorbs heat from surroundings
- Energy breaks ice crystal structure
- Endothermic ✓
The Correct Way
Think about energy flow:
- System absorbs heat → Endothermic (feels cold to touch)
- System releases heat → Exothermic (feels hot to touch)
Common examples:
Endothermic:
- Melting
- Evaporation
- Dissolving ammonium nitrate (cold pack)
- Photosynthesis
Exothermic:
- Freezing
- Condensation
- Combustion (fire)
- Neutralization reactions
How to Avoid
Ask: Where does the energy go?
- Into the system = Endothermic
- Out of the system = Exothermic
Sign convention:
- +ΔH = Endothermic (energy added)
- -ΔH = Exothermic (energy released)
Bonus Tips: General Study Habits
Before Exams
Don't:
- ❌ Memorize without understanding
- ❌ Cram the night before
- ❌ Skip practice problems
- ❌ Ignore sig figs and units
Do:
- ✅ Understand concepts first, then practice
- ✅ Space out study sessions
- ✅ Work through lots of problems
- ✅ Always include units and sig figs
During Exams
Check yourself:
- Did I balance the equation?
- Are my units consistent?
- Did I use Kelvin for gas laws?
- Do charges balance for ionic compounds?
- Is my answer reasonable? (If you get 10⁵⁰ moles of water in a cup, you did something wrong!)
Learning from Mistakes
When you get a problem wrong:
- Don't just look at the answer and move on
- Figure out WHERE you went wrong (concept? calculation? units?)
- Do a similar problem to make sure you've fixed the issue
- Write down the mistake and how to avoid it
Conclusion
These 10 mistakes account for the majority of lost points on chemistry exams:
- Confusing atoms, ions, and isotopes
- Not balancing equations
- Mixing up moles and grams
- Forgetting significant figures
- Ignoring charge balance
- Confusing molecular and empirical formulas
- Wrong oxidation states
- Using wrong units for gas laws
- Forgetting states of matter
- Mixing up endothermic and exothermic
The good news? Once you're aware of these pitfalls, they're easy to avoid. Slow down, double-check your work, and always ask yourself: "Does this make chemical sense?"
Pro tip: Use our interactive periodic table to check element properties, atomic masses, and oxidation states while studying!
Remember: Chemistry isn't about memorizing—it's about understanding patterns and thinking systematically. Master these common mistakes, and you'll see your grades improve dramatically.