Chapter 4: Compounds and Stoichiometry Flashcards
(26 cards)
Molecules:
Do ionic Compounds form molecules?
This is a combination of two or more atoms held together by covalent bonds.
• Are the smallest units of compounds.
• Made of two or more atoms of the same element
○ Like N2 and O2.
• Can also be made from two or more different atoms.
○ Like CO2, SOCL2, and more.
• Ionic compounds do not form true molecules because of the way in which the oppositely charged ions arrange themselves in the solid state.
Atomic Number:
This is the number on the top right corner of the box in the periodic table.
○ Example: Carbon has an atomic number of 6. Oxygen has a number of 8.
Atomic Weight:
This is the same as the atomic mass. This is the number or weight of the element itself. This number is below in the box in the periodic table.
○ Example: Carbon has an atomic weight of 12, oxygen has a weight of 16.
Molecular Weight:
Example: What is the molecular weight of SOCl2?
This is the sum of the atomic weights of all the atoms in a molecules.
○ Molecular weight of SOCl2:
S = 32 amu, O = 16 amu, Cl = 35 amu (2) = 70.
□ The molecular weight of SOCl2 is 32+16+70 = 118 amu.
Moles:
This is a quantity of any substance that equals to the number of particles that are found in an element.
To calculate the moles use the following equation:
§ Moles=(Mass of sample (g))/(Molar Mass (g/mol))
Molar Mass:
This is the mass of one mole of a compound that is expressed in g/mol.
Example: How many moles are in 9.53 grams of MgCl2?
MgCl2 = (24)+(2x35.5) = 95.3 g/mol # of moles = (9.53 g)/(95.3 g/mol)=0.10 moles
Equivalents:
This is how many moles of the thing we are interested in (protons, hydroxide ions, electrons, or ions) will one more of a given compound produce?
It includes gram equivalent weight, and normality.
Gram Equivalent Weight:
This weight is the amount of the compound, measured in grams, that produces one equivalent of the particle of interest.
§ Can be calculated by:
□ Gram equivalent weight=(Molar Mass)/n
® n is the # of particles of interest produced per molecule of the compound.
Example: How many grams would one need to produce one equivalent of hydrogen ions in H2CO3?
H2CO3 Molar Mass = 62, and there is 2 hydrogen which is the particle of interest so n = 2.
So (62 )/2=31 grams. So we would need 31 grams.
When the gram equivalent weight is known we can calculate how many equivalents?
When the gram equivalent value is known, we can see how many equivalents are there by using the equation:
Equivalents= (Mass of compound) / (Gram Equivalent Weigh)
Normality:
This is the measure of concentration, given in the units equivalents/L.
Molarity = Normality/n;
n = number of protons, hydroxide ions, electrons, or ions made or consumed by the solute.
Example: what is the normality of a 2M Mg(OH)2 solution?
The normality of a 2M Mg(OH)2 solution =
Molarity = Normality/n;
® n = 2, because there is 2 OH.
® The molarity is already given which is 2M.
® So Normality = Molarity x n = (2x2) = 4 N Mg(OH)2
Law of Constant Composition:
This law is when any pure sample of a given compound will contain the same elements in an identical mass ratio.
• Example: every sample of water will contain 2 hydrogen atoms for every one oxygen atom.
Empirical Formula:
This formula is the simplest whole-number ratio of the elements in the compound.
Molecular Formula:
This formula gives the exact number of atoms of each element in the compound and is a multiple of the empirical formula.
○ Example: For benzene the empirical formula is CH, but the molecular formula is C6H6.
For Ionic Compounds, what type of formula will they have?
Ionic compounds will only have empirical formulas.
Percent Composition:
This is the percent of the specific compound that is made up of a given element.
• To calculate this:
○ Percent Composition = (Mass of Element in Formula)/(Molar Mass)×100;
Example: What is the percent composition of chromium in K2Cr2O7?
Percent composition of chromium in K2Cr2O7:
□ Molar Mass of K2Cr2O7 = 294.2
□ Chromium = 52 and there are 2, so it will be 104.
□ Percent Composition = 104/294.2 = 0.3535 x 100 = 35.35 %
Combination Reactions:
This reaction is when two or more reactants combine to form one product.
• A + B –> C
• Example: when you burn hydrogen gas in air it forms water.
○ 2H2 + O2 –> 2H2O
Decomposition Reactions:
This reaction is when one reactant is chemically brown down into two or more products.
• A –> B + C
• Opposite to combination reaction.
• Chemically because of a result to heating, high-frequency radiation, or electrolysis.
• Example: breakdown of mercury.
○ 2HgO –> 2Hg +O2
Combustion Reactions:
This reaction is a special reaction that involves a fuel - usually a hydrocarbon -and an oxidant (normally) oxygen.
• Mostly forms two products of carbon dioxide and water.
• Example: Combustion of Methane
○ CH4 +2O2 –> CO2 +2H2O
Displacement Reactions:
This reaction is when one or more atoms or ions of one compound are replaced with one or more atoms or ions of another compound.
•Includes single and double displacement reactions.
Single-Displacement Reactions:
This reaction is when an atom or ion in a compound is replaced by an atom or ion of another element.
○ Example: Cu + AgNO3 –> Ag +CuNO3
