# SCH3U1 Grade 11 Trends, Nomenclature, Reactions Test

Topics:

1. Significant Digits
2. Trends of the Periodic Table
3. Types of Bonds
4. Properties of Ionic and Covalent Compounds
5. Lewis Structures
6. Nomenclature/Balancing Equations
7. Types of Reactions

Significant Digits

Rules:
1. all non-zero digits are significant
2. all leading and following zeros are not significant
3. all zeros between two other digits are significant
4. any zero that follows a digit and is to the right of the decimal point is significant

When performing multiplication or division, the value with the fewest number of significant digits is the number of significant digits used in the final answer

When performing addition or subtraction, the value with the fewest number of digits after the decimal place represents the number of decimal places in the final answer

Scientific Notation: used for numbers greater than 1000 or less than 0.001

To use, move the decimal to be placed after the first non-zero digit
Count the number of decimal places moved
If the decimal has moved to the left, the exponent is positive
If the decimal has moved to the right, the exponent is negative

Trends of the Periodic Table

– measured in picometers (1 X 10-12 m)
– distance from the nucleus to the valence orbital
– going right and up, the atomic radius decreases
– this is because going right, there is a greater force of attraction between the protons and electrons which makes a smaller radius
– going up the periodic table, there are less valence orbitals meaning the radius becomes smaller

Key Terms:

Shielding Effect: the effect of filled inner electron orbitals on the attractiveness of valence electrons to the nucleus

Effective Nuclear Charge: the apparent nuclear charge experienced by valence electrons, result of the shielding effect

Trend 2 Ionization Energy:

– amount of energy required to remove an electron
– going right and up on the periodic table will increase the ionization energy
– this is because fewer valence shells and more protons will make it require more energy to remove electrons

Trend 3 Electron Affinity:

– energy absorbed or released when an electron is added to a neutral atom
– going right and up the periodic table, the electron affinity will increase

Trend 4 Electronegativity:

– indicator of the relative ability of an atom to attract shared electrons
– smaller atomic radius usually means greater electronegativity
– going right and up the periodic table, electronegativity increases

Types of Bonds

Ionic:

– occurs between a metal and non-metal
– two ions bond to each other
– change in electronegativity is 1.7 – 3.3

Polar Covalent:

– occurs when non-metals share electrons
– has a positive and negative end
– change in electronegativity is 0.5 – 1.7

Covalent:

– occurs when non-metals share electrons
– does not have a positive/negative end
– change in electronegativity is 0 – 0.5

Properties of Ionic and Covalent Bonds

Ionic:

– high melting point
– does not conduct electricity as a solid
– conducts electricity in a solution
– high solubility in water
– solid state
– brittle, hard consistency as a solid

Covalent

– low melting point
– does not conduct electricity as a solid or in a solution
– not soluble in water
– non-polar compounds soluble in non-polar compounds
– polar compounds soluble in polar compounds
– can be solid, liquid, or gas
– soft, flexible, waxy consistency as a solid

Lewis Structures

– symbol of element is the center
– valence electrons are represented by dots surrounding the symbol of the element
– when dealing with ionic compounds, arrows are used to represent the loss/gain of electrons
– when dealing with covalent compounds, circles around electrons represent the sharing of electrons
– line structures simply use lines to show the sharing of electrons

Nomenclature/Balancing Equations

– will not be a major part of the test, so does not need to be reviewed too much

Types of Reactions

Synthesis:

– two elements combine to create a binary compound
– universal metal + element  compound
– multivalent metal + non-metal  compound
– non-metal + non-metal  molecular compound
– non-metal oxide + H2O will create an acid – will comprise of a hydrogen cation and a polyatomic anion containing oxygen
– e.g. SO2 + H2O  H2SO3(aq) (sulphic acid)
– metal oxide + H2O  metal hydroxide (base)
– e.g. Li2O + H2O  2LiOH(aq)

Decomposition:

– compound is broken down into simpler ones or elements
– binary compound decomposes into elements
– metal nitrate decomposing into metal nitrite + O2
– metal carbonate decomposing into a metal oxide + carbon dioxide

Combustion

– chemical reaction of a substance with oxygen producing one or more oxides, heat, and light
– inorganic combustion general formula: X + O2  XO
– organic combustion only occurs with hydrocarbons which are compounds composed of only hydrogen and carbon
– complete combustion of a hydrocarbon general formula: CxHy + O2  CO2 + H2O
– incomplete combustion of a hydrocarbon general formula: CxHy + O2  C(s) + CO + CO2 + H2O

Single Displacement

– chemical reaction in which one element in a compound is displaced by another element
– always refer to reactivity chart when dealing with single displacement reactions
– metal displacing another that is below it on the activity series in an ionic solution
– e.g. Cu + 2AgNO3  Cu(NO3)2 + 2Ag
– metal ion displacing hydrogen in an acidic solution
– e.g. Mg + 2HCl(aq)  MgCl2+ H2
– active metal + water  metal hydroxide + H2
– e.g. 2Rb + 2H2O  2RbOH + H2
– halogens displacing other halogens

Double Displacement

– chemical reaction in which positive ions of two ionic compounds exchange places, creating 2 new ionic compounds
– neutralization reaction: acid + base  water + “Salt” (ionic compound)
– e.g. H2SO4(aq) + Ca(OH)2(aq)  H2O + NaCl
– precipitation reaction: reaction with two ionic solutions that creates a precipitate
– e.g. 2LiCl(aq) + Pb(NO3)2  2LiNO3 + PbCl2