GEN CHEM Q1

SCIENCE

• Biological

• Physical

CHEMISTRY

→ Science that deals with the properties, composition and structure of substances

→ MAIN BRANCHES OF CHEMISTRY

• Organic Chemistry

• Inorganic Chemistry

• Physical

• Biochemistry

• Analytical Chemistry

PROPERTIES OF MATTER

Distinguished in 2 Types:

• Composition - different components

• Properties - qualities and attributes

Physical Property

• Identified using our senses and does not change in chemical composition

• Same molecule is present

• 2 TYPES OF PHYSICAL PROPERTY

1. Intensive — doesn’t depend on the amount 

Ex. density, color, conductivity, malleability (flattening), luster (shininess), ductility (stretchiness)

2. Extensive — amount or quality 

Ex. mass, volume, length

Chemical Property

• “Potential” to undergo chemical change or reaction by virtue of its composition 

• Ex. Corrosion of iron, flammability of paper

CLASSIFICATION OF SEPARATION TECHNIQUES

(1ST PPT)

MATTER

• Anything that has mass and occupies space

SOLID

• particles are packed tightly together (high density) 

• electrons constantly in motion, so the atoms have a small vibration, but they are fixed 

• very low kinetic energy 

• Definite shape, mass, volume 

• do not conform to the shape of the container in which they are placed

LIQUID

• particles are more loosely packed

• able to flow around each other

• indefinite shape, conform to the shape of the container

• incredibly difficult to compress 

GAS

• particles have a great deal of space between them

• high kinetic energy

• no definite shape or volume

PLASMA

• not a common state of matter here on Earth, but it may be the most common state of matter in the universe, according to the Jefferson Laboratory

• Stars — superheated ball of plasma

• consists of highly charged particles with extremely high kinetic energy

• The noble gasses are often used to make glowing signs by using electricity to ionize them to the plasma state. 

BOSE-EINSTEIN CONDENSATE

• created by scientists in 1995

• Eric Cornell and Carl Wieman - cooled a sample of rubidium to within a few degrees of absolute zero (at this low temp, molecular motion comes very close to stopping

• Since there is almost no kinetic energy being transferred from one atom to another, the atoms begin to clump together. There are no longer thousands of separate atoms, just one "super atom." 

CLASSIFICATION OF SEPARATION TECHNIQUES

• Physical State

• Composition

• Substances – composed of pure elements or chemically bonded elements

• Mixtures – composed of non-bonded substances

CHEMICAL SUBSTANCES

• a form of matter that has constant chemical composition and characteristic properties

• It cannot be separated into components without breaking chemical bond

• often called ‘pure’ 

Ex. Element (made up of a particular kind of atom and hence cannot be broken down)

Pure water (always has the same properties and the same ratio of hydrogen to oxygen whether it is isolated from a river or made in a laboratory)

diamond (carbon), gold, table salt (sodium chloride), and refined sugar (sucrose)

CHEMICAL MIXTURES

• material system made up of two or more different substances, which are mixed but not combined chemically

• identities of the individual substances are retained

• alloys, solutions, suspensions, and colloids

(2ND PPT)

PROPERTIES OF MATTER AND CHEMICAL FORMULA

DIFFERENT TYPES OF MATTER

COMPOSITION

— the different components of matter along with their relative proportions

PROPERTIES

— the qualities/attributes that distinguish one sample of matter from another

• Physical properties — observed by 5 senses, no changing of physical composition

• Intensive Physical Property

• density: ρ=m/v 

• color: pigment or shade

• conductivity: electricity to flow through the substance 

• malleability: if a substance can be flattened

• luster: how shiny the substance looks 

• ductility: the ability of substance to be stretched

• Extensive Physical Property

• mass: how much matter in the sample

• volume: how much space the sample takes up

• length: how long the sample is

• Chemical Property — its "potential" to undergo some chemical change or reaction by virtue of its composition

Ex. Corrosion of iron, Flammability of paper

CLASSIFICATION OF MATTER IN TERMS OF CHEMICAL COMPOSITION

PURE SUBSTANCE

– cannot be separated into other kinds of matter by any physical means

MIXTURE

–  physical combination of two or more substances in which each substance retains its own chemical identity

SEPARATION TECHNIQUES

FILTRATION

• process in which a filter medium is used to separate solid particles from liquid or gaseous mixture allows the liquid to pass thru (filtrate) but holds solid particles behind (residue) 

* insolubility - substance does not dissolve in a solvent

* solubility - solid (solute) + liquid (solvent) 

* miscibility - both liquid (solute and solvent) 

EVAPORATION

• technique used to separate out homogeneous mixtures that contain one or more dissolved solid drives off liquid component from solid ones typically involves heat

* evaporation - liquid to vapor (water boiling point: 100°C)

* vaporization - vapor to gas or liquid to gas 

DISTILLATION

• separate mixtures comprised of two or more liquid

• a purification process where liquid components are vaporized and then condensed and isolated

• The condensate that is collected is called “distillate”

• Types of water: alkaline, bottled, purified, distilled, mineral

CHROMATOGRAPHY

• Separation by passing it in a solution or suspension/ as vapor (gas in chromatography) through a medium (filter paper)

•  Separate and identify mixtures that are or can be colored

• Elements are the building blocks of compounds

• A compound is a substance composed of atoms of two or more elements chemically combined in a fixed proportion

CHEMICAL FORMULA

• symbolic expression of a compound or substance

• described as shorthand of expressing the types and the number of atoms in a substance

(3RD PPT)

ATOMIC STRUCTURE, FORMULAS, AND NAMES OF COMPOUNDS

CHEMICAL/MOLECULAR FORMULA

– representation of a molecule that uses chemical symbols to indicate the types of atoms followed by subscripts to show the number of atoms

STRUCTURAL FORMULA

– same as chemical but also shows how the atoms are connected in the molecule.

Example

1. Methane 

Chemical Formula: CH4

1 carbon, 4 hydrogen

2. Benzene

Chemical Formula: C6H6

NOMENCLATURE

• set of rules used for forming the names or terms in a particular field of arts or sciences

• purpose: to ensure that a spoken or written chemical name leaves no ambiguity concerning what chemical compound the name refers to

• most widely accepted format for nomenclature has been established by IUPAC

RULES FOR NAMING MOLECULAR COMPOUNDS: 

1. Remove the ending of the second element and add “ide” just like in ionic compounds. (sodium chlorIDE)

2. When naming molecular compounds prefixes are used to dictate the number of a given element present in the compound. (mono, bi/di, tri, tetra, penta,hexa, hepta, octo, nona, deca)

3. single first element  = no prefix. For example, CO is carbon monoxide, not monocarbon monoxide.

4. Extra vowels are removed.For example, one oxygen would not be monooxide, but instead it’s monoxide. 

EMPIRICAL FORMULA

– gives the simplest whole-number ratio of atoms in a compound

STEPS TO DETERMINE EMPI.FORM.

1. Assume a 100g sample of the compound so that the given percentages can be directly converted into grams.

2. Use each element's molar mass to convert the grams of each element to moles.

3. In order to find a whole number ratio, divide the moles of each element by whichever of the moles from step 2 is the smallest.

4. If all the moles at this point are whole numbers (or very close), the empirical formula can be written with the moles as the subscript of each element.

5. In some cases, one or more of the moles calculated in step 3 will not be whole numbers. Multiply each of the moles by the smallest whole number that will convert each into a whole number. Write the empirical formula.

Example 

A compound of iron and oxygen is analyzed and found to contain 69.94% iron and 30.06% oxygen. Find the empirical formula of the compound.

Given: Fe: 69.94%, O: 30.06%

Step 1: convert percent to grams

Fe → 69.94g

O → 30.06g

Step 2: convert to moles

Step 3:  divide both by the smallest result

Common Compounds 

1. Vinegar - solution composed of water and acetic acid

2. LPG (Liquefied Petroleum Gas) - composed of butane and propane, methane was replaced bcs highly flammable

3. MSG (Monosodium glutamate) - enhance umami flavor, na-ban dati, ok na now basta use moderately 

4. Mouthwash something 

5. Salt

1. Acetone - 2-propanone, colourless liquid used in the production of plastics and other industrial products

2. Sodium Chloride (NaCl)

3. Betadine - for minor wounds, can be mouthwash

4. Hydroperoxide - H2O2

5. Vinegar - solution composed of water and acetic acid

6. Aspirin - C9H8O4, reduce fever and relieve mild to moderate pain

7. Caffeine (?) - antioxidant, relieve mental and physical fatigue 

8. Mothballs - naphthalene 

9. Hydrochloric acid (HCl) - has two forms: aqueous solution (muriatic acid) and solid, pH level: 0.9-0.913???

10. Sucrose - food sweetener/preservative (glucose/dextrose - natural ba)

11. Sodium fluoride (NaF) - toothpaste

12. Calcium carbonate (CaCO3) - limestone

13. Solid carbon dioxide - dry ice

ATOMS

• “the basic building blocks of matter”. Atom is the basic of all matter. They are very small and consist of even tinier particles. Neutrons, Protons, and Electrons are the basic particles making up the atom. They join together with other atoms and create matter.

CHEMICAL FORMULA

• Empirical 

• Molecular

• Structural – H - O - H

• Molecular Geometry

Empirical formula -simplest formula! 

Nomenclature

Types of compound relative to bounding

1. Ionic  = Metal + Non-metal

2. Covalent = Non-metal + non-metal

ISOTOPES

• Are atoms of the same element (has the same no. of Protons but diff no. of Neutrons

Same # of protons but dif neutrons

• Radioactive decay occurs when protons exit the nucleus 

• When a neutron exits the nucleus, it turns into positive or negative

• Has the same chemical properties but medj diff in physical properties

Chemical Components

-googled

P - protons 

E- electrons

N - neutrons

Stable Isotopes 

• Any of two or more forms of an element whose nuclei contain the same no. of P and E but diff no. of N 

• Remain unchanged indefinitely 

Unstable (radioactive) Isotopes 

• Undergo spontaneous disintegration

Molecular formula 

• it gives the total no. of atoms of each element present in a molecule. 

STEPS

1. Determine the given

2. Get the empirical formula 

3. Multiply the no of atoms sa atomic mass niya tapos iadd lahat

4. Divide the given molar mass sa nakuhang sagot

5. Multiply mo yung empirical ans to the answer na nakuha mo sa step 4

Ex. The molar mass of a compound is 92g/mol. Analysis of a sample of the compound indicated that it contains 0.606g N and 1.390g O. Find its molecular formula. 

1. Given 

Molar mass: 92g/mol

0.606g N 

1.390g O 

2. Get the Empirical formula 

0.606g N/ 14 = 0.0432

1.390g O/ 16=0.0868

N- 0.0432/0.0432=1 

O-0.0868/0.0432=2

Empirical: NO2

3. Multiply 

N= 1 x 14=14

O=2 x 16=32

total= 48 

4. Divide 

92/48=1.917

Round off: 2

5. Empirical formula multiplied sa nakuhang sagot dito 

[NO2]2=N2O4

Atomic Structure, Formulas, and Names of Compounds 

Atomic Structure 

protons=atomic no.=electrons 

Atomic mass=protons+neutrons 

atomic mass = 16 

Atomic mass = p+n 

16=8+8

16=16

PERIODIC TRENDS – ewan ko ah pero ganito mga lumabas sa google

Melting point trends – Metals generally possess a high melting point. Most non-metals possess low melting points. The non-metal carbon possesses the highest melting point of all the elements. The semi-metal boron also possesses a high melting point.

Nuclear Fission - is a reaction in which the nucleus of an atom splits into two or more smaller nuclei. The fission process often produces gamma photons, and releases a very large amount of energy even by the energetic standards of radioactive decay. 

IONIC COMPOUNDS AND NAMING

Ionic Bonds 

• Ionic compounds - made when two ions join.

• Are due to an electrostatic attraction between ions of opposite charge. 

• Cations are positively charged ions (metals).

• Anions are negatively charged ions (non-metals).

CHEMICALS NAME SYMBOLS AND FORMULAS

Names are often based on a property or a discoverer. 

• Sodium - Arabic word “suda” means headache.

• Hafnium - discovered in Copenhagen. (Latin = Hafnia. 

• Lithium - Gk “lithos” means rock or stone

• Curium - for Marie Curie.

Symbols used in place of the name. 

• O - Oxygen

• N - Nitrogen

• C - Carbon

• H - Hydrogen

First letter is always capitalized in a symbol - other letters are always lower case: 

• Co -Cobalt

• CO -Carbon Monoxide

First letter is always capitalized in a symbol - other letters are always lower case 

• Sulfur - S

• Selenium - Se

• Strontium - S

• Sodium - Na

Formula - a combination of symbols showing composition. 

• Subscript - a small number that tells how many atoms of the element are in the compound and is written after the element it describes. 

• H2O means that two atoms of hydrogen are chemically bonded to one atom of oxygen. 

• If there is no number, “1” is implied.

Oxidation Number 

• The total number of electrons that an atom either gains or loses in order to form a chemical bond with another atom. The hypothetical charge of an atom if all of its bonds to other atoms were fully ionic.

Examples

• CaBr2 - calcium bromide 

• NaCl - sodium chloride 

• H2Se - hydrogen selenide

Formula writing 

Write the correct chemical formula that would form between 

a) calcium and fluorine 

b) Rubidium and oxygen 

c) Magnesium and phosphorus

NAMING COMPOUNDS

Positive ion (cation) is always named and written first!

Polyatomic Ions

• If multiples of the same polyatomic ion are used, then use parentheses to show the multiples.

Examples

• K2SO4 - potassium sulfate

• Li3PO4 - lithium phosphate

• BaCrO4 - barium chromate

• (NH4 )2Cr2O7 - ammonium dichromate

Stock Naming System 

• Used with transition elements and those other which can have multiple oxidation numbers (Fe +2,+3)? Pg. 222 

• Iron oxide could be Fe2O3 or FeO

• To clarify this, the oxidation number is stated immediately after the cation, then the anion is named as usual with its “-ide” ending in the case of binary compounds or the polyatomic name in the case of polyatomic ions. 

• Because iron has an oxidation number of two in the formula FeO, its name is iron two oxide. 

• But we use roman numerals to write it out. 

• So FeO is iron(II) oxide. 

• Remember the roman numeral refers to the oxidation number NOT the subscript!

MERCURY

• Mercury II is written Hg+2 

• Mercury I is written Hg2 +2 

• Why? 

• Hg+ + Hg+ = Hg2 +2

Covalently Bonded Binary Molecules 

• When naming things that bond covalently, (nonmetal to nonmetal), where oxidation numbers are not helpful, we use a system which uses prefixes. 

• Water, therefore, could be called dihydrogen monoxide. Di indicating there are two hydrogens and mono indicating one oxygen. (Note the other binary rules still apply giving oxygen the “-ide” ending. 

• What is the formula for trinitrogen pentoxide? 

• N3O5

QUANTITATIVE RELATIONSHIPS IN CHEMISTRY

Chemical Reaction 

• Is the change of a substance into a new one that has a different chemical identity. 

• Is usually accompanied by easily observed physical effects, such as the emission of heat and light, the formation of a precipitate, the evolution of gas, or a color change. 

• Absolute confirmation of a chemical change can only be validated by chemical analysis of the products!

TYPES OF CHEMICAL REACTION 

1. Combination or Synthesis Reaction

• two or more substances combine to form a new compound. 

2. Decomposition Reaction

• is a reaction in which a compound breaks down into two or more simpler substances.

3. Single-replacement Reaction

• is a reaction in which one element replaces a similar element in a compound. 

4. Double-replacement Reaction

• is a reaction in which the positive and negative ions of two ionic compounds exchange places to form two new compounds.

5. Combustion Reaction

• is a reaction in which a substance reacts with oxygen gas, releasing energy in the form of light and heat. Combustion reactions must involve O2 as one reactant. The combustion of hydrogen gas produces water vapor

Chemical Equation 

• A chemical equation consists of the chemical formulas of the reactants (on the left) and the products (on the right). The two are separated by an arrow symbol (“→” usually read aloud as “yields”). Each individual substance’s chemical formula is separated from others by a plus sign. As an example, the formula for the burning of methane can be written as follows: 

CH4 + O2 ➔ CO2 + 2H2O

• A balanced chemical equation has equal numbers of atoms for each element involved in the reaction are represented on the reactant and product sides. This is a requirement the equation must satisfy to be consistent with the law of conservation of matter

MOLE/MASS RATIO

• One of the fundamental laws of chemistry deals with the fact that we cannot (using chemical means) create or destroy matter. 

• When a reaction is run, the number of atoms of each specific type must be the same on both sides of the equation. 

• For some materials, it turns out that one element can combine with a second element in more than one ratio. Carrying out mass ratio calculations helped establish the law of multiple proportions.

Mole Ratio - is a conversion factor that relates the amounts in moles of any two substances in a chemical reaction. The numbers in a conversion factor come from the coefficients of the balanced chemical equation. The following six mole ratios can be written for the ammonia forming reaction.  N2 (g) + 3H2 (g) → 2NH3 (g)

Examples

1. How many moles of ammonia are produced if 4.20 moles of hydrogen are reacted with an excess of nitrogen? 

N2 (g) + 3H 2 (g) → 2NH3 (g)

2. How many moles of HCl will be produced when 249 g of AlCl3 are reacted according to this chemical equation? 

2 AlCl3 + 3 H2O(ℓ) → Al2O3 + 6 HCl(g)

3. How many grams of NaOH are required to fully consume 3.10 grams of H2SO4?

2NaOH (aq) + H2SO4 (aq) → 2H2O(l) + Na2SO4 (aq)

4. How many grams of NH3 will be produced when 33.9 mol of H2 are reacted according to this chemical equation? 

N2 (g) + 3 H2 (g) → 2 NH3 (g)

5. Determine the number of grams of SO 3 that can be produced by the reaction of 45.3 g of SO 2 and O2  

2SO 2 (g) + O 2 (g) → 2 SO 3 (g)

6. How many grams of nitrogen are in 25.0 g of (NH4 )2SO4 ? 

7. Lysine (C 6 H14 N 2 O 2 ) is an amino acid that is used in biosynthesis of protein. How many moles of carbon, hydrogen, nitrogen and oxygen are present in 0.75 moles of sample?

Avogadro’s Number - is the number of atoms, molecules, or other objects that makes up one mole of a substance. For example: 

6.022 x 1023 hydrogen atoms represent one mole of hydrogen. 

6.022 x 1023 water molecules represent one mole of water. 

6.022 x 1023 elephants represent one mole of elephants. 

THEORETICAL YIELD AND PERCENT YIELD

• To compute the percent yield, it is first necessary to determine how much of the product should be formed based on stoichiometry. This is called the theoretical yield, the maximum amount of product that could be formed from the given amounts of reactants. The actual yield is the amount of product that is formed when the reaction is carried out in the laboratory. The percent yield is the ratio of the actual yield to the theoretical yield, expressed as a percentage. 

Examples 

1. Potassium chlorate decomposes upon slight heating in the presence of a catalyst according to the reaction below: 

2KClO3 (s) → 2KCl (s) + 3O2 (g) 

In a certain experiment, 40.0g KClO3 is heated until it completely decomposes. What is the theoretical yield of oxygen gas? The experiment was performed, and the oxygen gas was collected, and its mass is found to be 14.9g. What is the percent yield for the reaction?

2. A worker reacts 30.5 g of Zn with nitric acid and evaporates the remaining water to obtain 65.2 g of Zn(NO3 )2 . What are the theoretical yield, the actual yield, and the percent yield?

Zn(s) + 2 HNO3 (aq) → Zn(NO3 )2 (aq) + H2 (g)

MASS RELATIONSHIPS IN CHEMICAL REACTIONS

Excess and Limiting Reagents

• When there is not enough of one reactant in a chemical reaction, the reaction stops abruptly. To figure out the amount of product produced, it must be determined which reactant will limit the chemical reaction (the limiting reagent) and which reactant is in excess (the excess reagent).

• One way of finding the limiting reagent is by calculating the amount of product that can be formed by each reactant; the one that produces less product is the limiting reagent.

Examples 

1. Consider this reaction:  

As(s) + O2 (g) → As2O3 (s) 

Suppose we start a reaction with 50.0 g of As and 50.0 g of O2 . Which one is the limiting reagent

2. A 5.00 g quantity of Rb are combined with 3.44 g of MgCl2 according to this chemical reaction: Rb(s) + MgCl2 (s) → Mg(s) + RbCl(s) 

a) Which one is the limiting reagent? 

b) What mass of Mg is formed, 

c) and what mass of what reactant is left over?

3. Calculate the mass of magnesium oxide possible if 2.40 g Mg reacts with 10.0 g O2 Mg+O2→MgO

4. Consider respiration, one of the most common chemical reactions on earth. 

C6H12O6 + O2 → CO2 + H2O + energy 

a) What reactant is in excess? And 

b) Calculate how much is left in the reaction of 25 grams of glucose and 40 grams of oxygen.

5. What is the greatest amount of aluminum chloride (in grams) that can be made with 114 g of aluminum and 186 g of chlorine? Which is the limiting reactant? Which is the excess? 

Al + Cl2 → AlCl3

GAS LAWS

Pressure - is defined as force/area. In chemistry, usually pressure comes from gases. The absence of pressure is called vacuum.

How do we Measure Pressure?

• Traditionally, pressure was measured using a barometer, which originally was just a tube with a column of liquid. The greater the height of the liquid, the higher the external pressure. Barometers can be used to measure pressure, which can be used to predict weather. 

UNITS OF PRESSURE

Gas Laws

• Created in the early 17th century, the gas laws have been around to assist scientists in finding volumes, amount, pressures and temperature when coming to matters of gas. The gas laws consist of three primary laws: Charles' Law, Boyle's Law and Avogadro's Law (all of which will later combine into the General Gas Equation and Ideal Gas Law).

1. Boyle's Law

• Robert Boyle discovered the correlation between Pressure P and Volume V (assuming Temperature T and Amount of Gas n remain constant):

2. Charles' Law

• French physicists Jacques Charles, discovered the correlation between Temperature in Kelvin T and Volume V (assuming Pressure P and Amount of Gas n remain constant):

3. Gay-Lussac’s Law

• The pressure exerted by a confined gas is directly related to the Kelvin temperature at constant volume. 

4. Avogadro's Law

• Amedeo Avogadro fixed Gay-Lussac's issue in finding the correlation between the Amount of gas(n) and Volume(V) (assuming Temperature(T) and Pressure(P) remain constant):

• Another form of the equation (assuming there are 2 sets of conditions, and setting both constants to each other) that might help solve problems is:

5. Combined Gas Law

• The law states that the ratio between pressure, volume, and absolute temperature of a gas equals some constant

6. Ideal Gas Law

• is the combination of the three simple gas laws. By setting all three laws directly or inversely proportional to Volume, you get:

• Next replacing the directly proportional to sign with a constant R you get:

7. Dalton's law of partial pressures

• States that the total pressure of a mixture of gases is equal to the sum of the partial pressures of the component gases:

• Dalton's law can also be expressed using the mole fraction of a gas, x:

STOICHIOMETRY

• Is a section of chemistry that involves using relationships between reactants and products in a chemical reaction to determine desired quantitative data

• Mole – Mole 

• Mole – Mass 

• Mass – Mass 

• Volume – Mole 

• Volume – Mass 

• Volume – Volume

• Avogadro’s Law states that ONE MOLE of ANY gas will occupy 22.4 L at STP (Standard Temperature and Pressure)

Examples: (walang sol sa ppt)

1. What volume of chlorine is produced when 45.0 g of aluminum chloride is decomposed? 

2AlCl3 → 2Al + 3Cl2

2. Find the mass of sugar (C6H12O6 ) required to produce 1.82 L of carbon dioxide gas at STP from the reaction described by the following equation: 

C6H12O6 + 6O2 → 6 H2O + 6 CO2

3. What volume of oxygen is used to convert 75.0 L of SO2 to SO3 ? Assume STP. 

2SO2 + O2 → 2SO3

4. How many ml of oxygen are required for the complete combustion of 47.2 g of butane, C4H10? *density of oxygen is 1.429 g/L 2C4H10 + 13O2 → 8CO2 + 10H2O

5. This reaction was carried out: 

CaCO3+ 2HCl→ CaCl2+ CO2+ H2O 

What would be the volume of CO2 (at STP) produced from the complete reaction of 10.0 grams of CaCO2 ?

Kinetic Molecular Theory 

• According to Kinetic Molecular Theory, gaseous particles are in a constant state of motion, moving at random speeds and in many different directions.

Five Basic Assumptions of the Kinetic Molecular Theory

• The volume occupied by the molecules of the gas is negligible compared to the volume of the gas itself. 

• No attractive and repulsive forces 

• Constant random motion 

• Kinetic energy is conserved 

• The average kinetic energy of the gas molecules is directly proportional to the absolute temperature

Graham's Laws of Diffusion and Effusion

Diffusion - is the gradual mixing of gases due to the motion of their component particles even in the absence of mechanical agitation such as stirring. The result is a gas mixture with uniform composition

Effusion - Is the escape of gaseous molecules through a small (usually microscopic) hole, such as a hole in a balloon, into an evacuated space

Graham’s Law - The rate of effusion of a gaseous substance is inversely proportional to the square root of its molar mass. This relationship is referred to as Graham’s law, after the Scottish chemist Thomas Graham (1805–1869). The ratio of the effusion rates of two gases is the square root of the inverse ratio of their molar masses:

Examples: 

1. Calculate the ratio of the rate of effusion of oxygen to the rate of effusion of hydrogen.

2. An unknown gas effuses 1.66 times more rapidly than CO2 . What is the molar mass of the unknown gas? Can you make a reasonable guess as to its identity?

3. It takes 243 s for 4.46 × 10-5 mol Xe to effuse through a tiny hole. Under the same conditions, how long will it take 4.46 × 10–5 mol Ne to effuse?

4. Under the same condition of temperature and pressure, how many times faster will hydrogen effuse compared to carbon dioxide?

5. Compare the rate of diffusion of carbon dioxide and ozone at the same temperature. 

6. A gas diffuses 0.25 times as fast at Helium. What is the molar mass of the unknown?

7. What is the rate of effusion for a gas that has a molar mass twice that of the gas that effuses at a rate of 3.62 mol/min?