CHEMISTRY SHENANIGANS

Today's class consisted of an activity to build molecules with the kits Mr. Doktor gave us. After dividing ourselves into groups of 2 or 3, the race begun.

In today's class, we created a smell in our Esterfication lab. My group chose

Here's a quick review if you do not remember how Esterfication works:

Amines:
end in -amine after the alkyl prefix
have a double bonded oxygen to the carbon chain and NH2 connected to the carbon chain

Name the following:


Answers:
1. ethanamide
2. methanamide



Amides:
end in -amide with the carbon prefix
have possibilities of primary (one carbon chain), secondary (two carbon chains), or tertiary (three carbon chains) chains connected to nitrogen

Name the following:




Answers:
1. dimethylamine
2. methylamine

ALDEHYDES
Naming: change the -e ending to -al

Ex. Draw Methanal (formaldehyde)


Ex. Benzaldehyde




CARBOXYLIC ACIDS
Naming: change the ending to -oic acid

Examples:


ESTERS
Naming: the primary chain has the -yl ending and the secondary chain takes the -oate ending

Examples:


ESTERFICATION
Always has H₂O

Check out this website for some help!
http://www.ausetute.com.au/esters.html

Halides:
-halogen atoms replace a hydrogen: Bromo, Chloro, Floro
EX. 2,2, dibromo 3 chloro, 1,3 difloro propane


Alchohols:
-OH or hydroxyl group
-change the ending to -ol
EX. 3 floro 2 methyl 4 pentaol


Ketones:
-oxygen atom doubled bonded to carbon
-change the ending -one
EX. 2 propaone (acetone)


Ethane:
-have an O joingin the two carbon chains together
-name carbon side chain with -yl ending and add 'ether'
EX. diethyl ether

ALKENES
- compounds with double bonds end in -ene
- a number in front of the parent chain tells where the double bond is
- more than 1 double bond changes the parent chain slightly

ALKYNES
- for compounds with triple bonds, use the -yne ending
- follow all the same alkene rules

Here's a quick recap:



And for some practice, check out this worksheet:
Organic Chemistry (with answers)

- there are more carbon compounds than all ionic compounds combined
- the study of carbon compounds is organic chem
- carbon can have multiple bonds and form many different shapes

Hydrocarbons have 3 types of formulas:
- molecular formulas (C6H14)
- condensed structural formula (CH3-CH2-CH2-CH2-CH2-CH3)
- structural formula


Nomenclature of Hydrocarbons
- One molecular formula can have a number of different structures
- Isomers are compounds that can be drawn in more than one way

How to name alkanes
1) Add "ane" to the longest chain with the correct suffix
2) Locate any branches by number carbon atoms (with the lowest numer)
3) Name branches with the appropriate suffix and -yl ending (Alkyl branches)
4) If there are more than one of the same alkyl group, number each one and add the multiplier number in front of the branch name

- the formation of a solution depends onthe ability of the solute to dissolve in the solvent
-solvation is the interaction btw solutes and solvents
-ionic solids (salts) are crystals made up of ions
-molecular solids are cystals made up of neutral molecules
-dissolving ionic solutions produces ions in a process called dissociation
-ionization is the break up of a neutral molecutle into charged particles

EX. CH3COOH--->CH3COO- + H+

-Determining concentrations is relatively easy

EX. What is the [Cl-] in a solution of 0.50M AgCl3
1)AgCl3--->Ag+ + 3Cl-
2)0.50 x 3 = 1.50M

- solvents and solutues can be polar or non-polar
- non-polar substances have equal charge distribution
- polar substances have unequal charge distribution
- in simpler terms, polarity depends on the symmetry of the structural diagram

Practice Questions: Determine whether each is polar or non-polar.
SiF4
SF4
PCl5
H2O
XeF4

Answers:
Non-polar
Polar
Non-polar
Polar
Non-Polar

Definition: bonds between molecules, there are 3 types

1. London Dispersion Force (L.D.F.)
- results from temporary electron diploes
- weakest intermolecular force
- increases as the number of electrons increases
- occurs in any compound that has electrons (aka EVERYTHING)

2. Dipole-Dipole Force
- results from a permanent dipole in molecules
- polar molecules experience this force
- polarity depends how much elements want electrons (electronegativity)
- the strength of the bond depends on the difference in electronegativity between the two atoms

3. Hydrogen Bonds (H-Bonds)
- this is a special type of dipole-dipole bond between H and O, F, or N.



Looking for a bit more information? Check out this site!

Want some practice? Check out this quiz!

- electrical condcution in solutions requires charged ions to be present
- ionic slutions dissociate when placed in water
NaCl(s)--H2O--> Na+(aq) + Cl-(aq)
- molecular solutions do not usually split into ions
C12H22O11(s)--H2O--> C12H22O11(aq)

Follow these steps to determine conductivity:
Is it a metal?
If yes, then its conductive. If no, is it a solid non-metal? If yes, then its non-conductive. If no, is it an acid or base? If yes, then its conductive. If no, is it ionic? If yes, then its conductive and if it's no, then its non-conductive.

Today's class was all about reviewing for the test on Friday. We were given an atomic theories review sheet and had to answer as many questions as possible and then check them over in class. As well, Mr. Doktor reminded us what we would need to know for the test:

1) Dalton, Thomson, Rutherford, Bohr models
2) Bohr energy level diagrams
3) The differences between ions, atoms, isotopes
4) Atomic Structure (atomic number, atomic mass, number of electrons, etc)
5) Trends on the Periodic Table (mass, charge, ionization energy, size, reactivity)
6) Metals, non-metals, metalloid properties
7) Lewis dot diagrams
8) Structural diagrams

G'luck!

In today's class, we separated ourselves into groups and used a conductivity sensor to determine the conductivity of 9 substances.

Here are my results:

SOLUTION | CONDUCTIVITY | IONIC OR MOLECULAR? | ACID OR BASE?
Acetic Acid | 1518μs/cm | molecular | acid
Hydrochloric Acid | 16650μs/cm | molecular | acid
Hydrogen Sulfate | 32350μs/cm | ionic | acid
Copper Chloric Acid | 4156μs/cm | ionic | acid
Sodium Chloride | 4167μs/cm | ionic | base
Ammonia | 1500μs/cm | molecular | base
Sodium Hydroxide | 3300μs/cm | ionic | base
Sucrose | 0μs/cm | molecular | acid
Ethyl Alcohol | 0μs/cm | molecular | base

Take a look at this video about conductivity and pH:

Covalent Bonds
-electrons are shared between non-metals
-to draw Lewis Dot Diagram:
1. Add the valence e- in all atoms
2. Identify which atom can form the most number of bonds. This will be the central atom
3. Bonds between two atoms are represented by a e-. This represents 2e-
4. Any e- not creating bonds are placed in pairs around the remaining atoms
5. All valences levels must be filled, all electrons must be used

Double and Triple Bonds
- some compounds from more than one bond between two elements

http://www.youtube.com/watch?v=tOpke6cpqWY&feature=pyv&ad=4006919948&kw=covalent%20bonds
http://www.youtube.com/watch?v=mAjrnZ-znkY

Electronegativity
- atoms affinity for electrons
- electronegativity increase from left to right and from bottom to top

-atoms are electically neutral
-# of protons = # of electrons
-ions have different # of protons and electrons
-ions cane either be positive (lost e) or negative (extra e)
-cation = positive ion
-anion = negative ion

EX. Determine how many electrons each of the ions have. What type of ion are they(cation/anion)?
-Ca2+ lost 2e, cation, 18e
-Ag+ lost 1e, cation, 46e

EX. Determine how many protons,neutrons,electrons the following substances have.
-76As3- p=33 n=43 e=36
-201Au+ p=79 n=122 e=78

Bohr Diagrams for Ions
-energy level bohr diagram
8e <--- Ca+ 8e 2e 40Ca 20 Chemical Bonds
-a bond is an electrostatic attraction between particles
-bonds occur as elements try to achieve noble gas electron configuration
---> noble gases (usually) dont form compounds/bonds
---> in noble gasese the outermost energy level have stable octets
-metals lose electrons (oxidize)
-non-metals gain electrons (reduced)

Lewis Dot structure
-atoms can be repersented by dot diagrams
---> dots represent electrons
---> only valence electrons shown
-write the atomic symbol for the atom
---> this represents the nucleus and filled inner electrons levels
-one dodt is used to represent outer energy level elctrons
---> one e is placed in each orbital before any pairing occurs
---> beginning with the 5th e, pairing can occur up to max. of 8e

EX.

Ionic Bond
-electrons are transtered from metal to non-metal
---> no dots are shown on metal
-"charged" species is written in brackets

EX.

Ion Charge: the electrical charge that forms on an atom when it gains or loses electrons
- elements on the left side of the Periodic Table tend to form positive ions (cations)
- elements on the right side of the Periodic Table tend to form negative ions (anions)
- elements are arranged in columns or families by their similar ion charge(s)

• from the left side, elements begin with a positive charge that increases down the right of the period until it reaches a non-metal
• from the right side, elements begin with a negative charge that increases negatively to the left of the period until it reaches a metal
• the transition metals which have been positioned in the center (indicated on the chart by a solid checkered line) because most of them have more than one charge

Chemical Reactivity: the rate at which a chemical substance tends to undergo a chemical reaction time

- for metals, chemical reactivity increases as you move towards the left and down
- for non-metals, chemical reactivity increases as you move towards the right and up

ex. Potassium is more reactive than Beryllium because it is farther to the left and farther down. (Potassium also has much more electrons and is only one electron away from reaching a full valence shell.)


Ionization Energy: the energy needed to remove electrons from an atom; also referred to as ionization potential

- an atom with layers of inner electrons has stronger ionization energy because the inner electrons protect the outer electrons from the force of attraction the protons give off

ex.Chlorine has stronger ionization energy than Lithium

Vertical columns are the groups or chemical families: Alkali Metals, Alkaline Earth Metals, Transition Metals, Halogens and Noble Gases
- Hydrogen is its own group
- elements in the same chemical family have similar physical and chemical properties.



Alkali Metals
- in group 1
- highly reactive and reactivity increases as you go down
- have only one electron in their outer shell and are lose that electron in ionic bonding with other elements
- react with non-metal
- usually have lower densities than other metals
- malleable, ductile, good conductors of heat and electricity
- have low melting points, below 200 °C
- soft and can be cut with a knife

Alkaline Earth Metals
- 2nd group
- have two electrons in their outer shell
- have low electronegativities
- less reactive than Alkali Metals but they will burn in air if heated. T
- react with water
- shiny

Transition Metals
- are the 38 elements in groups 3-12 of the periodic table
- are very hard
- have high melting and boiling points
- low ionization energies
- high electrical conductivity
- are malleable, they can be shaped and bent.

The Halogens
- in group 17 of the periodic table
- are highly reactive non-metals with strong and unpleasant odours
- will burn flesh and do not react well with water
- fluorine and chlorine are gases at room temperature, bromine is a liquid and iodine and astatine are solids.

Noble Gases
- group 18
- most stable and unreactive elements
- colourless, odourless gases at room temperature
- have high ionization energies and low boiling points

- Dmitri Mendeleev started organizing every known element and is credited with creating the first real periodic table of elements
- he sawa pattern and left spaces in the table for elements that were undiscovered

- there are 118 elements on the Periodic Table
- 7 periods and 18 columns
- metals on the left and non-metals are on the right
- elements with similar properties are in the same vertical column
- Atomic Number: Number of protons of the nucleus of each atom of an element
- Atomic Mass: Mass of an average atom of an element (tends to increase along with atomic number)
- Ion charge: Electric charge that forms on an atom when it gains or loses electrons

Mr. Doktor wasn't in class today and we had a sub. Instead of reviewing homework, we were assigned our group chemistry projects and started working on them.

Just a quick recap:

This is how two isotopes of hydrogen would look like. In essence, most of their properties are retained except for certain structural details.

-each element gives of a specific colour of light
-these are known as emission spectra, unique to each element
-If electrons absorb energy they can bumped to a higer level
-when the fall to a lower level, they release energy as light

Atomic Structure:
-atoms are made up parts called subatomic particles:
1)protons- positive, mass 1, located in nucleus
2)neutrons- neutral, mass 1, located in nucleus
3)electrons- negative, mass 1/1837, located outside



Atomic Number:
-atomic number is the number of protons



Isotopes:
-the number of protons dertermine they type of element
-changing number of neutrons chages isotopes of the element
-all isotopes have the same chemical properites


Mass Number:
-mass number is the total number of protons and neutrons
-give symbol A
-different isotopes have different masses
-mass number = atomic number + number of neutrons
-A = Z+N

- atoms are electrically neutral
-two different models can be used to describe electron configuration
i) Energy Level Model
ii) Bohr Model
-Electrons occupy or orbitals
i) 2 e in the first orbital
ii) 8 e in the second orbital
iii) 8 e in the third orbital
- these are called octet

Hybridized Orbitals
- the first of the Bohr Levels is the 1s-orbital and it holds 2e-
- the second level contains the 2s, 2px, 2py, 2pz orbitals. They combine (hybridize) to from one 2sp3 orbital

In today's class, Mr. Doktor discussed atomic theory. There were seven significant people who's ideas contributed to the development of atomic theory: the Greeks, Lavoisier (late 1700s), Proust(1799), Dalton (early 1800s), J.J. Thomson (1850s), Rutherford (1905), and Bohr (1920s).

Some important points..

The Greeks:
- in 3000 BC, Democritus said that atoms were indivisible particles
- first mention of atoms
- only a conceptual model
- was the accepted view for over 2000 years

Lavoisier:
- Law of conservation of mass
- Law of definite proportions
*wasn't a true atomic theory because it did not discuss what atoms were or how they were arranged

Proust:
- if a compound is broken down into its constituents, the products exist in the same ratio as in the compounded form
- experimentally proved Lavoisier's Laws

Dalton:
- atoms are solid, indestructible spheres (like billiard balls)
- provides for different elements


J. J. Thompson:
- used the raisin bun model
-solid, positive spheres with negative particles embedded in them
- first atomic theory to have positive (protons) and negative (electron) charges
- introduced the idea of nucleus

Rutherford:
- showed that atoms have a positive, dense center with electrons outside it
- resulted in a planetary model
- explains why electrons spin around nucleus
- suggests atoms are mostly empty space
- should be unstable (electrons and protons should attract and destroy the atom)

Bohr:- electrons must only exist in specific orbitals around nucleus
- explains how valence electrons are involved in bonding
- explains the difference between ionic and covalent bonding
- resolves the problem of atomic instability
- includes the neutron (discovered in 1932)
- explains atomic emission spectra

Review Questions:
1. Who first mentioned atoms? Democritus
2. Who introduced the idea of nucleus? J. J. Thomson
3. Who stated that atoms were indestructible? Dalton
4. Who explained bonding and 'levels' around the nucleus? Bohr
5. Who explained why electrons spin around the nucleus? Rutherford

In today's class, we were put to the test as Mr. Doktor challenged us to create 2.00g of Barium Chloride as a precipitate for a grand prize of... donuts. Yes that's right, even better than good grades!! Haha. Anyways, we had to create our own procedure and were given different chemical compounds to choose from in order to create Barium Chloride for ourselves. Now that's it over and done with, the only question is: who gets the donuts?

Mass to Mass EXAMPLES:
Lead (IV) nitrate reactions with 5.0 g of Potassium iodide.
Grams of Lead (IV) Nitrate are required?
Pb(NO3) + 4 KI -----> 4 KO3 + PbI45.0 g x 1 mol KI/165.9 g x 1 Pb(NO3)4/4KI x 455.2/ 1 mol Pb(NO3)4 = 3.4 gif a 100 mL solution of 2.0 M H2SO4 is neutralized by sodium hydroxide what mass of water is produced?
H2(SO4) + NaOH -----> HOH + Na2(SO4)2.0 mol/L x .1 L = .200 mol H2(SO4) x 2 HOH/1 H2(SO4) x 18.0 g/mol = 7.2 g H2OPercent Yield

The theoretical yield of the reaction is the expected *calculated* amount. The experimental amount is the actual yield% yield = actual/theoretical x 100

IE: He production of urea CO(NH2) is given by:
2NH3 + Co2 -----> CO(NH2) + H2O. 47.7 g of urea are produced when 1 mole of CO2 reacts, find the actual yield, theoretical yield and percent yield.

The actual yield is 46.7 g.
The theoretical yield: 1 mol CO2 x 1 CO(NH2)2/1 CO2 x 60.1 g/ 1 mol = 60.1 g
Percent Yield = 47.7 g/60.1 g x 100 = 79.4%

-usually one reactant gets used up first
-one reactant is completely consumed
-this determines how much product is produced
-'Guess' which reactant is limiting, and check how much of it is required

What is thelimiting reactant when 125g of P4 reacts with 323g of Cl2 to form phosphorous tricholoride?

1. P4+6Cl2--->4PCl3
2.125g X 1mol/124g X 6mol of Cl2/1mol P4 X 71g/mol = 431g of Cl2
3. Cl2 is the limiting reactant

Deternmine the theoretical yield of the previous equation

323g X 1mol Cl2/71g X 4mol PCl3/6mol CL2 X 137.4g/1mol PCl3 = 417g

Here's a link for more info on limiting reactants:
http://www.chem.tamu.edu/class/majors/tutorialnotefiles/limiting.htm

In today's class, we did a lab. Last class, we had to complete the prelab which involved our prediction for the outcome of the experiment.

The Experimental Design: 2.00 g of Strontium nitrate is dissolved in 50mL of water and then reacted with excess Copper (II) sulphate (3.00g). The product is a precipitate (Strontium sulphate and Copper (II) nitrate. After mixing the solutions, the precipitate will be seprated by filtration, dried and weighed.

My group's prediction results:

1. What is the balanced chemical equation for this reaction?
Sr(NO₃)₂ + CuSO₄ --> SrSO₄ + Cu(NO₃)₂

2. If 2.00 g of Strontium nitrate completely reacts, how many grams of Strontium sulphate should be produced?
2.00 g Sr(NO₃)₂ x (1 mol/211.6g) x (1 mol/SrSO₄/1 mol Sr(NO₃)₂) x 183.7g/mol = 1.74 g



We completed the first part of the lab (dissolving the chemicals, forming a precipitate, and seprating it through filtration) and the second half (drying and weighing) will occur next class.

Ex. When 2.0g of oxygen reacts with nitrogen monoxide, how many moles of nitrogen dioxide are produced?
O2+2NO >>> 2NO
2.0gx (1 mol of O2/32.0g) x (2mol NO2/1mol O2) = 0.125 mol of NO2

Ex. If 3.0g of O2 with nitrogen monoxide, what volume of nitrogen dioxide is produced at STP?
1. O2 + 2NO2 >>> 2NO2
2. 3.0g x (1mol O2/ 32.0g) x (2mol NO2/1mol O2) x (22.4L/mol) = 4.2L

Ex. In the formation of Copper(II) Oxide, 4.0g of copper react. How many moles of copper oxide are produced?
1.2Cu + O2 >>> 2CuO
2. 4.0g x (2 Cu/63.5g) x (2mol of CuO/ 2mol Cu) = 0.126 mol CuO

Ex. What mass of water is produced when 5.5L of hydrogen is burnt(reacted w/O2)?
1. 2H2 + O2 >>> 2H2O
2. 5.5L x ( mol/22.4) x (18.0g/mol) = 4.4 g

Ex. Lead (IV) Nitrate reacts with 2.0g of potassium iodide. How many moles of leader (IV) nitrate are needed?
1. Pb (NO3)4 + 4KI >>> PbI4+ 4KNO3
2. 2.0g x (1mol PbI4/ 714.8g) x (1 mol Pb(NO3)4/ 1mol PbI4) = 2.8x 10-3 mol

-Coefficients in balanced equations represent moles
-Also conversion factors

Ex. (In class) If a 0.15 mole sample of methane reacts with oxygen. How many moles of each product are produced?
1)CH4+2O2--->CO2+2H20
2)0.15molCH4 x 1molCO2/1molCH4 = 0.15 mol CO2
3)0.15molCH4 x 2molH2O/1molCH4= 0.30 mol H20

-Always what you need/what you know
-Converting to mass requires one addtional step

Ex. (In class)How many grams of Bouxite are required to produce 3.5 mol of Al?
1) 3.5molAl x 2molAl2O3/4molAl = 1.75 mol Al2O3
2) 1.75mol x102g/1mol=178.5g

Today, we went over the Molar Enthalpy and Combustion worksheet and did a lab. In our lab, we placed water in a can and lit a candle underneath for several minutes. Afterwards, with the measurements we took, we tried to make an educated guess of how much energy was released or absorbed in the experiment. It looked a lot like this:



Questions from the homework:

1. A car engine burns 250g of octane (C8H18). If 1.19x10^3 kJ of heat are produced, what is the molar enthalpy of the combustion of Octane?
Answer: 250g ÷ 114g/mol (of Octane) = 2.19 mol
1.19x10^3 kJ ÷ 2.19 mol = 543 kJ/mol

4. When ethanol C2H5OH is burned, heat is released. For every mole of ethanol burned, 1300 kJ of heat are released.
a) If 2.5g of ethanol are burnt, how much heat will be released?
Answer: 2.5g x mol/46g x 1300kJ/mol = 70.65 -> 71 kJ

Questions for practice:

ex. What is the molar enthalpy of CO2 (g) in the reaction for the burning of butane below?
2 C4H10 (l) + 13 O2 (g) 8 CO2 (g) + 10 H2O (g) H = -5315 kJ
Answer: Molar enthalpy, change in H = 5315 kJ ÷ 8 mol = 664 kJ / mol.

Calorimetry
To measure heat absorbed/released by water we need to know THREE important things

1) Temperature change ( 'C)
2) Amount of water (g,kg, mL, L)
3) Specific heat capacity (KJ/ Kg 'C)

deltaH= mCdeltaT

m= mass of water
C= specific heat capacity
deltaT= change of temperature *liquid water is 4.19 KJ/Kg '


MOLAR ENTHALPY
Heat obsorbed/ realeased by one mole

When a candle (C H ) is burnt, heat is released according to reaction"
25 52

C H + 380 ---> 25CO + 26H O+ 11000 KJ
25 52 2 2 2

ex. 9.0 grams of charcoal were completely consumed in a bomb calorimeter. If 2.0 L of water absorbed all of the heat released by the charcoal, and if the temperature of the water increased from 20.25 to 56.04oC, what is the molar enthalpy of carbon?
Answer: deltaH= mCdeltaT
= 2000g(4.19 J/(g C))(56.04-20.25)
= 299 920 J
dH = - C = - 299 920 J = -299.92 kJ
n = 9.0g / 12.01 g/mole = 0.7493 mole
DH/n = -299.92 kJ/ 0.7493 mole = -400.27kJ/mole of C = - 4.0 X 102 kJ/mole

- Rxns that release heat are exothermic (hot)
- Rxns that absorb heat are endothermic (cold)

Heat is a form of energy
All chemicals have energy stored in them. Stored chemical energy is called enthalpy.
Enthalpy is stored chemical energy.

Enthalpy of gasoline > Enthalpy of water

Exothermic rxns convert enthalpy into heat.
2C8H18 + 25O2 >>> 16CO2 + 18 H2O + Heat
high enthalpy low enthalpy

Enthalpy has a symbol H
Change in enthalpy is Delta H (triangle H)

Exothermic

• delta H is negative
• products are lower than reactants
• heat is released

Endothermic

• delta H is positive
• reactants are lower than products
• heat is absorbed

Delta H and Moles
coefficients can stand for moles or molecules

1)Synthesis
2Al+3F2--->2AlF3

2)Decomposition
*always assume the products are elements*
4H3PO4--->6H2+P4+8O2

3)Single Replacement
*compounds must always have a metal and a nonmetal*
3Mg+2Al(NO3)3--->2Al+3Mg(NO3)2

4)Double Replacement
MgCl2+K2SO4--->MgSO4+2KCl

5)Neutralization
*Between acids and bases*
*always forms H2O*
H2SO4+2KOH--->2H2O+K2SO4

6)Combustion (burning)
2Mg+O--->2MgO --Metallic
CH4+2O2--->CO2+2H20 --Hydro carbon

In today's class, we discussed how to balance equations involving carbon, hydrogen, and oxygen. We learned that in combustion reactions, the products will always be carbon dioxide and water.

ex. (in-class) CH₄ + 0₂ -> CO₂ + H₂O
becomes CH₄ + 20₂ -> CO₂ + 2H₂O

Also, we found out that we could use fractions while balancing equations. Because the number in front of the atom or compound says how many moles of it there is, it is possible to have a fraction as there can be 1/2 or 7/2 of a mole.

ex. (in-class) C₈H₁₈ + O₂ -> CO₂ + 9H₂O
becomes C₈H₁₈ + 25/2 O₂ -> 8 CO₂ + 9H₂O

Briefly, we discussed alcohols.
CH₃OH - methanol
C₂H₅O₅ - ethanol

In addition, we were given a number of acids we should know.
Hydrochloric acid (HCl)
Nitric acid (HNO₃)
Sulfuric acid (H₂SO₄)
Phosphoric Acid (H₃PO₄)
Acetic Acid (CH₃COOH)

Try some for yourself!
1. CO₂ + H₂ -> CH₄ + H₂O
2. C₂H₆ + O₂ -> CO₂ + H₂O

Answers:
1. CO₂ + 4H₂ -> CH₄ + 2H₂O
2. 2C₂H₆ + 7O₂ -> 4CO₂ + 6H₂O

Need to some more help or examples? Check out the videos below: