General Chemistry · Unit 01

Intro to Chemistry: Significant Figures, SI Units, and Lab Safety

Intro to chemistry sets up the foundation for the whole course: what chemistry studies, how to use common lab equipment correctly, how to measure with SI units and significant figures, how to stay safe in the lab, and how the scientific method works — before you move on to matter and the rest of the course.

What you'll learn

Explain what chemistry is. Identify common laboratory equipment and choose the right tool for the job. Measure and record data using SI units and significant figures. Read a graduated cylinder and use an electronic balance correctly. Apply lab safety rules before working with chemicals, glassware, or heat. Identify and understand the steps of the scientific method.

1.1 Start Here: What Chemistry Is Actually About

Chemistry is the study of matter, and the change it undergoes. Basically, that means chemists want to know what a substance is, how much of it is there, what happens when it changes, and how energy is involved.

These are the same big questions that keep showing up all year, even when the unit topic changes:

Chemistry question	What you are trying to figure out
What substance is present?	Identify the material or particle involved.
How much is present?	Measure mass, volume, amount, or concentration correctly.
What changed?	Track how matter or energy changed during a process.
What do the measurements show?	Use data to support a conclusion.

▸

Why this unit matters

This unit gives you the measuring, math, and lab habits that make later chemistry feel much less overwhelming. You can also come back here any time you need a reset on the basics.
You can also learn and practice reading tools like balances for mass and graduated cylinders for liquid volume so those skills feel normal before they show up inside harder problems.
Significant figures, SI units, and careful setup come back in atomic structure, moles, stoichiometry, thermochemistry, gas laws, solutions, equilibrium, and acids and bases.

1.2 Lab Equipment: Which Tool to Use and Why

Teachers typically have students learn common lab equipment in chemistry classes. Lots of equipment may seem similar to others (e.g. beakers, graduated cylinders, test tubes, flasks), but it is important to know which one to use for a specific use or situation:

Task	Best tool	Why
Hold or mix a liquid	Beaker	Good for holding and mixing, not for precise volume
Measure liquid volume accurately	Graduated cylinder	Better for measured volume than a beaker
Prepare one exact final volume	Volumetric flask	Built for one exact marked volume
Deliver a precise variable volume	Burette	Used when the delivered amount must be measured carefully
Transfer an exact volume	Pipette	Designed to measure and move a precise amount
Measure mass	Electronic balance	Reads mass directly; tare removes the container mass
Heat and stir a solution	Hot plate / stirrer	Designed for controlled heating and mixing
Support glassware over a heat source	Ring stand + clamps	Holds equipment in place during heating
Transfer solids	Spatula / scoopula	Moves solids cleanly without using your hands
Rinse glassware	Wash bottle	Lets you direct distilled water where you need it

▸

Quick review

If the question asks for an accurate measured volume, choose a graduated cylinder, pipette, or burette, not a beaker.
If the question asks for sample mass only, place the weigh boat on the balance first and press TARE before adding the sample using a scoopula or mini-spatula.
Important: "holds liquid" and "measures liquid" are not the same job. Beakers hold liquid but do not measure it accurately like a graduated cylinder, pipette, or burette.

1.3 Measurement and SI Units: Writing Data to Use in the Correct Equations and Keep Track of It

Science uses the International System of Units (SI) so measurements mean the same thing everywhere. Important: every measurement has a number and a unit. A number by itself is not enough. Calculations in other topics like atomic structure, moles, stoichiometry, thermochemistry, gas laws, solutions, equilibrium, and acids and bases will only work if your units are under control. Learn about SI Units in chemistry here:

▸

Every measurement is a number and a unit together. One without the other is not a measurement.
Write both the number and the unit every single time. You will lose points on exams, quizzes and homework if you don't.
Keep units visible through every step of the calculation. Do not take shortcuts. It always pays off.

SI Base Units

Quantity	Unit	Symbol
Length	meter	m
Mass	kilogram	kg
Time	second	s
Temperature	kelvin	K
Amount of substance	mole	mol
Electric current	ampere	A
Luminous intensity	candela	cd

Common Metric Prefixes

Prefix	Symbol	Multiplier	Example
mega-	M	106	1 Mm = 1,000,000 m
kilo-	k	103	1 km = 1,000 m
deci-	d	10-1	1 dm = 0.1 m
centi-	c	10-2	1 cm = 0.01 m
milli-	m	10-3	1 mm = 0.001 m
micro-	µ	10-6	1 µg = 0.000001 g
nano-	n	10-9	1 nm = 10-9 m

Temperature in chemistry is almost always in degrees Celsius for everyday lab measurements. You'll need to convert back and forth to Kelvin when you get to gas laws and thermodynamics. That's where the simple temperature equation, Kelvin = degrees Celciuis + 273 is used.

Temperature Conversions

K = °C + 273.15

°C = (°F − 32) × 59

°F = (°C × 95 ) + 32

▸

Chemists use Celsius (°C) for everyday lab measurements and Kelvin (K) for gas law and thermodynamics calculations.
Many intro chemistry problems round 273.15 to 273. Follow the directions given.
0 K (absolute zero) is the coldest possible temperature — there are no negative Kelvin values.

1.4 Significant Figures: Which Digits Count and Which Do Not

Sig figs are kind of weird at first, because most people do not run into them in everyday life and have to wonder, "How many sig figs do I need here?" One place you might see them is at a gas station. A gallon of gas might be posted as 2.99 9/10 dollars. That really means 2.999 dollars. Those three decimal places matter when the price gets calculated at the pump.

Another place sig figs show up is calories on a food label. For example, a label might say the serving size is 28 grams and there are 180 Calories in that amount. Why not 182 Calories or 177 Calories? It is because the ones place is not a sig fig. Only the tens place is significant on that label. If Calories were measured at 186, the label would round to 190.

Significant figures communicate the precision of a measurement. The big question is simple: which digits came from the measurement, and which zeros are only placeholders? If you are confused, fix it now. Sig figs keep showing up in calculations all the way through moles and stoichiometry.

Analog instrument: record all certain digits plus one estimated digit.
Digital instrument: record every digit shown on the display.

Rules for Counting Sig Figs

Rule	Example	Sig figs
All non-zero digits are significant	4,523	4
Zeros between non-zero digits are significant	4,023	4
Leading zeros are NOT significant	0.0047	2
Trailing zeros with a decimal point ARE significant	2.500	4
Trailing zeros without a decimal are NOT significant	2500	2
Scientific notation removes all ambiguity	2.50 × 103	3

How to Type Scientific Notation In chemistry, 9.53 × 1023 means 9.53 followed by 23 powers of ten.
In answer boxes, type it with e notation: 9.53e23.
For small numbers, use a negative exponent: 4.56 × 10-4 is typed as 4.56e-4.

Rewrite in Scientific Notation

6,020,000 = 6.02 × 106 because the decimal moves 6 places to the left.

0.000456 = 4.56 × 10-4 because the decimal moves 4 places to the right.

When you type them in an answer box, use 6.02e6 and 4.56e-4.

Quick Zero Test

Leading zeros only locate the decimal, so they do not count.

Middle zeros between non-zero digits do count.

Trailing zeros count only when the number format shows measured precision, such as a decimal point or scientific notation.

Sig Figs in Calculations

This is where the two rules get mixed up. The operation type — not the number of digits — tells you which rule applies.

Multiplication/Division: Answer has the same number of sig figs as the measurement with the fewest sig figs.
Addition/Subtraction: Answer is rounded to the same decimal place as the measurement with the fewest decimal places.
Exact numbers (counts, definitions) have unlimited sig figs and do not limit your answer.

Examples 4.52 × 2.1 = 9.492 → rounds to 9.5 (2 sig figs, limited by 2.1)
12.36 + 1.2 = 13.56 → rounds to 13.6 (tenths place, limited by 1.2)

▸

Common mistake

Very important: Do NOT round intermediate steps.
Carry extra digits through the entire calculation and round only the final answer.
Important:When you need to show that trailing zeros ARE significant, use a decimal point (2500.) or scientific notation (2.500 × 103).
Do not confuse decimal-place rules with sig-fig rules. Addition and subtraction care about decimal places. Multiplication and division care about sig figs.

1.5 Dimensional Analysis: How to Convert Units and Set Up Chemistry Problems

Want to get better at chemistry test and quiz problems? Here are two habits that will help a lot: show your work every time and use dimensional analysis, also called the factor-label method. If you always write out your work and use units to guide you, you will make fewer mistakes and get a lot more confident with chemistry problems.

Factor-Label Routine

1. Start with the given number and given unit.

2. Multiply by a conversion fraction with the starting unit on the bottom and the wanted unit on the top.

3. Use the correct ratio, such as 10 mm/1 cm or 1 L/1000 mL, so the unit you do not want cancels.

4. Keep multiplying by new conversion factors until the only unit left is the unit for your answer.

Important Idea

A conversion factor is a ratio equal to 1, so it changes the unit without changing the amount.

Example: 10 mm/1 cm and 1 cm/10 mm are both correct ratios. Choose the one that makes the unwanted unit cancel.

Do the unit cancellation first, then multiply the numbers, then round the final answer.

Worked Example 1

Convert mL to L

Problem: Convert 250.0 mL to liters.

Given: 250.0 mL

Find: L

Step 1 — Identify the starting unit and the wanted unit

Start with 250.0 mL. The answer needs to be in L.

Step 2 — Put mL on the bottom so it cancels

250.0 mL × 1 L/1000 mL = 0.2500 L

mL cancel, leaving L ✓.

Step 3 — Keep the unit you wanted and round at the end

250.0 mL = 0.2500 L. The conversion factor is exact, so the answer keeps the 4 significant figures from 250.0 mL.

Worked Example 2

Use Density to Find Volume

Problem: A sample has a mass of 13.5 g. Its density is 2.70 g/mL. Find the sample volume.

Given: 13.5 g; 2.70 g/mL

Find: mL

Step 1 — Choose the ratio that cancels grams

The given unit is g, but the answer needs mL. Since density = g/mL, flip it so g is on the bottom:

1 mL / 2.70 g

Step 2 — Multiply so the grams cancel

13.5 g × 1 mL/2.70 g = 5.00 mL

g cancel, leaving mL ✓.

Step 3 — Check that the final unit matches the question

V = 5.00 mL. This works because the density ratio was arranged to remove grams and leave milliliters.

Worked Example 3

Use Density and Then Convert to Kilograms

Problem: A gold sample has a volume of 25.0 mL. Gold has a density of 19.3 g/mL. Find the mass of the sample in kilograms.

Given: 25.0 mL; 19.3 g/mL

Find: kg

Step 1 — Start with the relationship that connects volume to mass

Density tells you grams per milliliter, so start by multiplying the volume by the density to get grams.

mass = volume × density

Step 2 — Use density to change mL into g

25.0 mL × 19.3 g/1 mL = 482.5 g

mL cancel, leaving g ✓.

Step 3 — Convert grams into kilograms

482.5 g × 1 kg/1000 g = 0.4825 kg

g cancel, leaving kg ✓.

Step 4 — Round the final answer to the correct significant figures

m = 0.483 kg. The given numbers 25.0 mL and 19.3 g/mL each have 3 significant figures, so the final answer should also have 3 significant figures.

Worked Example 4

Convert nm to m

Problem: Convert 8.50 nm to meters.

Given: 8.50 nm

Find: m

Step 1 — Use the metric prefix relationship

1 nm = 10-9 m, so a matching factor-label ratio is 1 m / 109 nm.

Step 2 — Put nm on the bottom so it cancels

8.50 nm × 1 m/109 nm = 8.50 × 10-9 m

nm cancel, leaving m ✓.

Step 3 — Write the final answer with the correct unit

8.50 nm = 8.50 × 10-9 m. The unit conversion is exact, so the answer keeps 3 significant figures.

▸

Quick check before you move on

If the wrong unit is still showing, the conversion factor is upside down.
For multi-step problems, do not stop until the unit left is the unit the question asked for.
Keep units written through the entire setup. They tell you whether the math makes sense.

1.6 How to Read a Graduated Cylinder Correctly

A graduated cylinder measures liquid volume, the amount of space the liquid takes up. If your graduated cylinder is glass, it will almost always form a curved surface at the top called the meniscus. For water and most aqueous solutions, that curve goes downward. It is extremely important to read the liquid at the meniscus from the bottom of that curve, at eye level. Many students read from the top of the curve or look from the wrong angle — both give the wrong number. It ruins the reading, and you will get the problem wrong or do your lab calculations incorrectly.

Start Here: Reading Routine

1. Find the smallest marked interval.

2. Read the bottom of the meniscus at eye level.

3. Record one estimated digit beyond the smallest marking.

10 mL Cylinder

Read the bottom of the meniscus.
This reads 6.8 mL.

Estimated between 6 & 7 mL graduations

100 mL Cylinder

This sketch shows the meniscus position clearly.
The recorded reading is 42.0 mL.

The trailing zero matters because the final digit is estimated

25 mL Cylinder

Estimate to ±0.1 mL.
This reads 18.5 mL.

Halfway between 18 & 19 mL lines

▸

Precision rule: Read the bottom of the meniscus for most aqueous solutions. Record one estimated digit beyond the smallest marked interval.

1.7 How to Use an Electronic Balance Without Ruining the Measurement

An electronic balance measures mass in grams. In chemistry, that tells you how much of a substance you have for a reaction or how much was produced in a reaction. It is really important to understand the TARE button. TARE zeros out the balance after you place a container on it, like a weigh boat or a beaker. That way, the container does not count toward the mass of your sample. The usual process is: put the container on the balance, press TARE, and then add your sample. Now the balance will only show the mass of the substance you added. This is one of the most common lab mistakes, but it is also one of the easiest to avoid.

Electronic Balance

Display reads 15.432 g. The TARE button zeros out the mass of a container.

Correct Balance Technique

Power onAllow the balance to warm up briefly.

Place containerPut your weigh boat or beaker on the pan.

Press TAREThe display zeros out so the container mass is removed.

Add sampleCarefully add your substance. Record all digits shown.

Record & cleanWrite the full reading. Remove the sample and clean the pan.

▸

Never skip taring.
TARE zeros the display so only the sample mass shows.
Always tare with the container already on the pan, before you add anything.

Problem	Result
Forgot to tare	The display includes container + sample.
Did not record all shown digits	You lose precision.
Removed the tared container and kept going	The reading is no longer valid.

Do Not Miss These Common Errors

Parallax: Reading a graduated cylinder from above or below eye level gives an incorrect meniscus position.
Forgetting to tare: Always tare the balance with the container before adding your sample.
Rounding too early: Carry all digits through calculations; round only the final answer.
Confusing mass and weight: Mass (kg, g) is a property of matter; weight is the gravitational force on that mass.
Hypothesis ≠ prediction: A hypothesis must be testable and falsifiable. "I think plants will grow better" is not a hypothesis.
Forgetting units: A number without a unit is not a measurement.

1.8 Lab Safety: What You Need to Do Before Anything Goes Wrong

Lab safety is very important and keeps you and your lab mates safe and healthy. Start here: match each hazard to the right action before you touch chemicals, glassware, or heat. That is how you protect yourself, your lab partner, and your data.

Personal Protective Equipment (PPE)

🥽

Safety goggles — worn whenever chemicals, heat, or glassware are in use. Regular glasses do not count.

🧤

Gloves — nitrile gloves protect skin from corrosive, toxic, or skin-absorbed chemicals. Remove before touching your face.

🥼

Lab apron/coat — protects clothing and skin from spills. Long pants and closed-toe shoes are required.

Before lab

These three happen before you touch anything. If you skip them, the rest of the list doesn't protect you.

Read the entire procedure before beginning.
Know the location of the eyewash station, safety shower, fire extinguisher, and first aid kit.
Put on goggles and any required protective clothing before handling chemicals or glassware.

During lab

Notice: most of these are about habits, not knowledge. They only protect you if they become automatic.

Never eat, drink, or apply cosmetics in the lab.
Read labels carefully before using a chemical.
Never smell a chemical directly — waft the vapors toward you with your hand.
Add acid to water, not water to acid.
Keep long hair tied back and loose items away from flames and chemicals.

If something goes wrong

Report spills, breakage, and exposure to your teacher immediately.
Use safety equipment right away when needed.
Dispose of chemicals only as instructed.

Safety Response Routine

Flammable near a burner → move it away from the flame and keep the cap closed.

Corrosive on skin → rinse immediately and tell the instructor at once.

Toxic vapor warning → avoid direct inhalation and follow the lab ventilation instructions.

1.9 How Chemists Test an Idea and Defend a Claim

Science is a way of learning about the world using evidence and experiments. It helps us answer questions we can actually test by asking, testing, measuring, and using data to support our conclusions.

Step 1: Ask a testable question

Write a question you can answer with observations or measurements. Example: "Does higher water temperature make an antacid tablet dissolve faster?"

Step 2: Use what is already known

Use what you already know to make a smarter test. For the antacid example, you might already know that higher temperature makes particles move faster. That kind of background knowledge helps turn your idea into a real hypothesis instead of just a guess.

Step 3: Write a real hypothesis

State what you predict and why. A hypothesis is testable. It is often written in if-then form.

Step 4: Set up a fair test

Choose one variable to change. Choose what to measure. Keep the other conditions the same. Include a comparison condition when needed.

Variable type	Definition	In the example
Independent	What you deliberately change	Water temperature
Dependent	What you measure as a result	Time for the tablet to dissolve
Constants	Everything else held the same	Tablet brand, water volume, container, stirring
Comparison	Condition used for a fair check	Room-temperature water

Step 5: Collect and organize the data

Record everything clearly as it happens — not from memory afterward.

Write observations and measurements the moment you make them.
Be specific: "the solution turned pale yellow" is useful. "It changed" is not.
Organize results so you can actually compare them in the next step.

Step 6: Make the conclusion match the evidence

Look at what the data actually show, not what you hoped they would show. Good science follows the evidence.

If the data support the hypothesis, say so with specific evidence.
If they do not, that is still a real result — report it honestly.
Conclusions are based on evidence, not opinion.

Step 7: Share the work so others can check it

Other scientists need to be able to repeat your experiment. If they cannot, the result does not hold up. Write the method clearly enough that someone else could repeat your experiment.

▸

Common mix-up: theory vs. law

A scientific law describes what happens consistently.
A scientific theory explains why or how it happens.
Neither word means "guess" in science. A theory is not a weak law, and a law does not explain the cause.

Use these like a teacher-guided check, not like a guessing game. Commit to an answer first, then compare your reasoning. That is how you catch weak spots before they become habits. For more reps, use the Unit 01 Practice page.

Read the Graduated Cylinder

Read This First

Commit to a reading before clicking Check. Read the bottom of the meniscus and estimate one digit beyond the smallest marked interval.

Scenario: Your lab partner has filled a graduated cylinder with a liquid reagent. You need to confirm the exact volume before it gets added to the reaction.

Cylinder size

Choose the best reading

Feedback

Choose a reading above, then click Check.

Do not reveal the answer until you commit to a reading.

Use the Balance Correctly

Read This First

Build a valid sample-only mass reading before you check.

The balance only gives a usable sample mass after the container is on the pan and the display has been tared to zero.

Scenario: A student needs the mass of a powder sample using a weigh boat. Work through the procedure, then check whether the displayed mass is actually usable.

Place the container on the pan to begin.

// Action log

Work through the procedure, then click Check.

Count the Significant Figures

Read This First

Decide which digits count before you reveal the rule.

Ignore placeholder zeros first. Then decide whether trailing zeros count in this form of the number.

Scenario: Read one measurement, choose how many significant figures it has, then commit before reveal.

Choose the Rounding Rule

Read This First

Choose the rule that limits the final answer before you calculate.

Ask whether the operation is addition/subtraction or multiplication/division. Then decide whether decimal places or sig figs control the answer.

Scenario: Read one calculation setup, choose the rule that applies, then commit before reveal.

Example 1 · Applying the Scientific Method

Problem: A student wants to test whether higher water temperature changes how fast an antacid tablet dissolves.

Given: The student can change water temperature and measure dissolving time while keeping tablet brand, water volume, container, and stirring method the same.

Find: The independent variable, dependent variable, and what result would support the hypothesis.

Step 1 — Write the question

Question: Does higher water temperature make an antacid tablet dissolve faster?

Step 2 — Hypothesis

If the water temperature increases, then the tablet will dissolve faster, because particle motion increases at higher temperature.

Step 3 — Identify the variables

Independent variable: water temperature. Dependent variable: time to dissolve. Constants: tablet brand, water volume, container, and stirring method.

Step 4 — Describe the procedure

Run trials at different temperatures, measure how long the tablet takes to dissolve, and compare the results.

Step 5 — Decide what would support the hypothesis

If higher-temperature trials dissolve faster than lower-temperature trials, the data support the hypothesis.

Example 2 · Significant Figures in a Multiplication

Problem: A strip of metal foil is measured as 3.21 cm long and 1.5 cm wide. Calculate the area and report it with the correct number of significant figures.

Given: length = 3.21 cm, width = 1.5 cm

Find: area with the correct number of significant figures

Step 1 — Write the formula

area = length × width

Step 2 — Identify sig figs in each measurement

3.21 cm → 3 sig figs | 1.5 cm → 2 sig figs

Step 3 — Calculate (keep extra digits for now)

3.21 × 1.5 = 4.815 cm2

Step 4 — Round to correct sig figs

The limiting measurement is 1.5 cm (2 sig figs). Round 4.815 to 2 sig figs:

area = 4.8 cm2

Common wrong answers: 4.815 cm2 means you forgot to round. 4.80 cm2 keeps too many significant figures.

Example 3 · Reading a Graduated Cylinder

Problem: A 50 mL graduated cylinder has major markings every 10 mL and minor markings every 1 mL. The bottom of the meniscus sits halfway between the 27 mL and 28 mL lines. What is the volume?

Given: smallest marked interval = 1 mL; meniscus bottom sits halfway between 27 and 28 mL

Find: the recorded volume with correct precision

Step 1 — Identify the graduation interval

Minor markings are every 1 mL, so you can estimate to the nearest 0.1 mL.

Step 2 — Read the bottom of the meniscus

The meniscus bottom is halfway between 27 and 28 mL → estimated position = 27.5 mL.

Step 3 — Record with correct precision

V = 27.5 mL

Three sig figs — the 2 and 7 are certain; the 5 is estimated.

Example 4 · Using the TARE Function

Problem: A beaker has a mass of 87.450 g. A student places it on the balance, presses TARE, then adds salt until the display reads 14.230 g. What is the mass of the salt? What would the balance show if the student forgot to tare?

Given: beaker mass = 87.450 g; tared reading after salt is added = 14.230 g

Find: the salt mass with TARE and the display reading if TARE were skipped

Step 1 — With TARE pressed first

After taring, the balance display = 0.000 g (beaker mass removed). Adding the salt gives:

mass of salt = 14.230 g

Step 2 — Without TARE

If TARE was skipped, the display would show:

87.450 g + 14.230 g = 101.680 g

This reading includes the beaker mass, so it is not a valid sample-only mass.

Step 3 — Sig figs on the answer

The key decision is whether the display shows the sample only or the container plus sample. On a digital balance, record all digits shown on the display.

Example 5 · Unit Conversion to Scientific Notation

Problem: How many seconds are in 5.00 years? Use 365 days/year and report the final answer in scientific notation.

Given: 5.00 y; 365 d = 1 y; 24 h = 1 d; 60 min = 1 h; 60 s = 1 min

Find: seconds in scientific notation

Step 1 — Set up one conversion chain so the units cancel

5.00 y × 365 d/1 y × 24 h/1 d × 60 min/1 h × 60 s/1 min

y cancel → d cancel → h cancel → min cancel → s ✓ remains

Step 2 — Multiply after the units cancel

5.00 × 365 × 24 × 60 × 60 = 157,680,000 s

The counted conversion factors are exact here, so the 5.00 years controls the significant figures.

Step 3 — Rewrite the final answer in scientific notation

Final Answer

1.58 × 108 s

Type as 1.58e8.

Move the decimal 8 places to the left, then keep 3 significant figures because 5.00 has 3 significant figures.

→

Best way to lock in Unit 01

After the Unit 01 Practice page, use the broader ChemUnlocked Practice Hub when you want mixed review, then read How to Study Smarter, Not Longer if you need a better plan for spaced review and correcting mistakes.

General Chemistry · Unit 01 · Intro to Chemistry & Lab Safety