Test #4- Chapter 2 sec3 + chapter 3 (all sections) - Abbabbo

Accuracy

How close a measured value is to the accepted (true) value

Precision

How close repeated measurements are to each other

Difference between accuracy and precision

Accuracy is closeness to the true value; precision is consistency of measurements

Percent error formula

∣experimental−accepted∣÷accepted×100

Percent error

A measure of how far an experimental value is from the accepted value

Significant figures

Digits in a measurement that are known with certainty plus one estimated digit

Purpose of significant figures

To show the precision of a measurement

Rule: sig figs in nonzero digits

All nonzero digits are significant

Rule: sig figs in zeros between numbers

Zeros between nonzero digits are significant

Rule: sig figs in leading zeros

Leading zeros are NOT significant

Rule: sig figs in trailing zeros with a decimal

Trailing zeros ARE significant if a decimal point is present

Rule: sig figs in trailing zeros without a decimal

Trailing zeros are NOT significant unless specified

Sig figs in multiplication/division

Answer has the same number of sig figs as the value with the fewest sig figs

Sig figs in addition/subtraction

Answer has the same number of decimal places as the value with the fewest decimals

Scientific notation

A way to express numbers as a product of a number between 1 and 10 and a power of 10

Scientific notation format

a×10^n, where 1 ≤ a < 10

Convert to scientific notation

Move the decimal so one nonzero digit remains on the left

Expand scientific notation

Move the decimal right or left based on the exponent

Direct relationship (graph)

A straight line with positive slope where y increases as x increases

Inverse relationship (graph)

A curved graph where y decreases as x increases

Direct variation equation

y = kx

Inverse variation equation

y = k/x

Constant (k) in a direct relationship

k = y ÷ x

Constant (k) in an inverse relationship

k = xy

Philosophic idea

Based on reasoning and logic, not experimentation

Scientific theory

A well-tested explanation based on repeated experimentation and evidence

Democritus' view of matter

Matter is made of tiny, indivisible particles called atoms

Aristotle's view of matter

Matter is continuous and infinitely divisible

Law of conservation of mass

Mass is neither created nor destroyed in a chemical reaction

Applying conservation of mass

Total mass of reactants equals total mass of products

Law of definite proportions

A compound always contains the same elements in the same fixed ratio by mass

Law of multiple proportions

When two elements form more than one compound, the masses combine in simple whole-number ratios

Dalton's Atomic Theory (5 points)

All matter is made of atoms; Atoms of the same element are identical; Atoms of different elements differ; Atoms combine in whole-number ratios; Chemical reactions rearrange atoms

Modern atomic theory difference #1

Atoms are divisible into subatomic particles

Modern atomic theory difference #2

Atoms of the same element can have different masses (isotopes)

Cathode ray tube experiment conclusion

Atoms contain negatively charged particles (electrons)

Electron

A negatively charged subatomic particle with very small mass

Three contributions of J.J. Thomson

Discovered the electron; Determined electron charge-to-mass ratio; Proposed the plum pudding model

Plum pudding model

Electrons embedded in a positively charged sphere

Rutherford gold foil experiment

Alpha particles fired at thin gold foil

Why gold foil was used

Gold can be hammered into extremely thin sheets

Key observation of gold foil experiment

Most particles passed through; some were deflected

Conclusion of gold foil experiment

Atom is mostly empty space with a dense, positive nucleus

Nucleus

Small, dense, positively charged center of the atom

Millikan oil drop experiment

Measured the charge of an electron

What Millikan determined

The electron's charge (and later mass)

Proton

Positively charged particle in the nucleus

Neutron

Neutral particle in the nucleus

Location of electrons

Electron cloud surrounding the nucleus

Relative charge of proton

+1

Relative charge of neutron

0

Relative charge of electron

−1

Relative mass of proton

1 amu

Relative mass of neutron

1 amu

Relative mass of electron

~0 amu (1/1836 of a proton)

Mass ranking (greatest to least)

Neutron ≈ Proton > Electron

Nuclear forces

Strong forces that hold protons and neutrons together in the nucleus

Atom vs nucleus size

The atom is much larger than the nucleus

Atomic number

Number of protons in an atom

Mass number

Number of protons + neutrons

Isotope

Atoms of the same element with different numbers of neutrons

Average atomic mass

Weighted average of an element's isotopes based on abundance

Location of atomic number on periodic table

Whole number, usually above the element symbol

Location of average atomic mass

Decimal number below the element symbol

Protons in an atom

Equal to the atomic number

Electrons in a neutral atom

Equal to the number of protons

Neutrons in an atom

Mass number − atomic number

Nuclide

A specific isotope of an element

Relative mass

Atomic masses are compared to carbon-12

Carbon-12 definition

Assigned a mass of exactly 12 amu

Atomic mass unit (amu)

One-twelfth the mass of a carbon-12 atom

Average atomic mass calculation

Sum of (isotope mass × percent abundance as a decimal)

Mole

A counting unit equal to 6.02 × 10²³ particles

Avogadro's number

6.02 × 10²³ particles per mole

Molar mass

Mass of one mole of a substance in grams

Molar mass location

Same numerical value as average atomic mass on the periodic table

Convert moles to grams

Multiply by molar mass

Convert grams to moles

Divide by molar mass

Convert atoms to moles

Divide by Avogadro's number

Convert moles to atoms

Multiply by Avogadro's number

Atoms → grams (2-step)

Atoms → moles → grams

Grams → atoms (2-step)

Grams → moles → atoms