Chem fundamentals (unit 1)

Matter

Matter

Physical material; anything with mass that takes up space

Pure substances

Matter that has a definite composition, one that does not change, and has distinct properties. They can only be separated by chemical reactions.

Elements

Pure substances that cannot be decomposed into simpler substances.

Atoms

The smallest building block of matter. Each element is composed of 1 type of atom.

Ex: C

Molecules

2+ atoms, can be same or different.

ex: O2, H2O

Compounds

Pure substances composed of 2+ different elements. They can only be separated chemically.

Ex: H2O

Physical properties

Can be observed without changing the identity or composition of matter. They are the result of IMFAs between structures.

ex: melting point and color

Intensive properties

Properties that are independent of quantity.

ex: boiling point, odor

Extensive properties

Properties that are dependent on quantity.

Ex: mass, volume

Chemical properties

Observed by destroying substance, they result from chemical reactions.

Mixtures

A combo of 2+ pure substances. Each substance maintains its own properties. Mixtures can be separated into its pure substances.

Homogeneous mixtures

Mixtures that are uniform throughout. The components are evenly distributed. They look pure but they aren’t since they aren’t chemically combined

Solution

Homogeneous mixture with small particles that don't scatter light.

ex: brass, copper sulfate (aq)

Colloid

Homogeneous mixture with large particles that scatter light.

ex: milk

Heterogeneous mixtures (suspensions)

Mixtures that aren't evenly distributed, you can see layers/each part

ex: wood, granite, rice pudding

Physical change

Changes physical appearance, not composition.

ex: ice to water is a state change

Chemical change

Substance changes into a different substance.

ex: heat of combustion, flammability

Distillation

Process that depends on the boiling points to separate mixtures.

ex: boiling NaCl and water. Water evaporates, leaving NaCl behind

Chromatography

Process that depends on the differing size and polarity of substances to separate mixtures.

Filtration

Process of pouring a mix of solids and liquid through filter paper to separate them. The liquid passes through, solid stays behind

ex: coffee

Dalton’s atomic theory:

1. Each element is composed of extremely small particles (atoms)

2. All atoms of a given element are identical to each other

   1. All O2 atoms are the same, all N2 atoms are the same

3. Atoms of 1 element can’t be changed into atoms of different elements by chemical reactions.

   1. O2 can’t turn into N2

4. Compounds are formed when atoms of more than 1 element chemically combine.

   1. O + N (elements)→NO (compound)

Law of conservation of mass:

matter isn’t created or destroyed, just rearranged

Law of constant composition/definite proportion:

given compounds always have same elements in the same proportion. The ratios are fixed

ex: water is always H2O, a 2:1 ratio of H to O

Law of multiple proportions:

compounds with different ratios of the same atoms are different

ex: H2O2 is different from H2O

Democritus:

made first atomic model in 400 BC

proposed that all matter is made up of atoms (small, solid, indivisible particles)

Model: ball

Dalton:

determined that each element is made up of atoms, created atomic theory

model: ball

Thomson:

through cathode ray tube experiments, determined that there are negatively charged electrons

because electrons contribute a small fractions of atom’s mass, they are small

model: plum pudding

Rutherford:

discovered protons and nucleus

Most of the atom’s mass comes from dense + nucleus and most of the volume is empty space (electron cloud)

model: nuclear

Chadwick:

through nuclear bombardment, he found the neutron

Beta radiation

high speed electrons with charge -1

Alpha radiation

charge: +2

Gamma radiation

high energy, no mass, no charge

Nucleus:

contains protons and neutrons with an overall + charge.

very small and dense (1x10^-15 m)

Electron cloud:

contains negatively charged electrons

almost no mass but most of atom volume (1-5 x10^-10 m)

Angstrom:

1x10^-10m=100pm

Atomic mass unit (amu)= __g

1.66054x10^-24 g

Proton:

1.0073amu (1amu)

+1

Electron:

5.486x10^-4amu (0amu)

-1,

Neutron:

1.0087amu (1amu)

0

Isotopic Notation

A= mass #

Z=atomic #

q=charge

mass number

protons+neutrons

Isotopes:

same # protons but different # neutrons, differing mass numbers

atomic number

just protons, used to identify element

Charge in isotope notation

#protons vs electrons

1amu=____g

1.66054x10^-24 g

1g=____amu

6.02x10^23 amu

How to calculate atomic mass

the number from the periodic table is dependent on isotopic abundance

Σ (isotope mass)(isotope abundance)

Ex: 9/16 atoms have a mass of 70, 6/16 have a mass of 72, 1/16 have mass of 74

AM=70(9/16)+72(6/16)+74(1/16)=71amu

Spectrometer

1. Get atoms into gas phase and convert them into ions (cations)

2. When gas phase cations made, they’re accelerated towards negative grid

3. Only a narrow beam of ions can pass

4. Beam passes through magnet poles that deflect ions

5. Ions separated into their masses (isotopes)

Mass spectroscopy:

uses spectrometer to determine the mass of an element/molecule

Provides mass of ions and relative abundance, allows us to calculate atomic mass

Periods:

the rows

Groups:

the columns

Molecular vs empirical formula

molecule is the actual # of atoms in a molecule while empirical is the smallest ratio

Molecular: H2O2

Empirical: HO

Formula weight

the sum of each atomic weight in a substance

FW of H2O: 2(1.008)+1(16)=18.016 g

% composition

the mass contributed by each element

% comp= #atoms of element (atomic weight) /formula weight x100

How to get the empirical formula from % of each element

Base the calculation on 100.g of compound. It’s easier

Determine # moles of each element for 100.g of compound

Divide each mole value by the smallest mole value to get the ratio

Multiply by the integer to get a whole number formula

How to get molecular formula from empirical

whole # multiple=molecular weight/empirical formula weight

Then multiply empirical formula by that multiple

Combustion analysis to get empirical formula

1. Use mass of CO2 to find the amount of C in organic substance

2. Use mass of H2O to find amount of H in organic substance

3. If there’s oxygen in the organic, subtract Cmass and Hmass to get Omass by itself

4. Once you have masses of each element, proceed like before, get mole substance ratios

Coefficients:

the relative # of molecules in a reaction

• Ex: 2H2 +O2 →2H2O shows 2 molecules of H2 reacting with 1 molecule of O2 to form 2 molecules H2O

Limiting Reactant

Reactant that limits how much product is formed. Once LR runs out, the reaction stops

Theoretical yield:

quantity of product calculated to form (100% completion, no error)

Experimental yield

Quantity of product that actually formed in lab

Percent yield

How much product you got compared to the true amount

experimental/theoretical x100 = ___%

Percent error

How far off you were from the theoretical value

exp-theo/theo x 100 = ± ___%

Molarity

Concentration of moles/L

Dilution

Adding water to make a concentration lower

C₁V₁=C₂V₂

Titration

combining solution with unknown concentration w reagent of known concentration

Equivalence point:

where stoich equivalent quantities are brought together

Indicator

dye that changes color as passing equivalence point