Chemistry Midterm

Unit 1

1. Matter: Matter is anything that has mass and occupies space. Solid, liquid, or gas.

2. Difference between an Observation and an Assumption:

Observation: An observation is the process of gathering information using your senses or instruments. It is based on factual data.

Assumption: An assumption is a belief or judgment made without direct evidence. It's an idea you accept without proof.

3. Difference between Qualitative and Quantitative Observations:

Qualitative Observation: Describes qualities or characteristics that cannot be measured numerically (e.g., color, texture, smell).

Quantitative Observation: Involves measurements or quantities, expressed in numbers (e.g., length, mass, temperature).

4. Difference between Intrinsic and Extrinsic Properties:

• Intrinsic Properties: Properties that are inherent to the substance and do not depend on the amount or size of the material (e.g., density, boiling point, color).

• Extrinsic Properties: Properties that depend on the amount or size of the material (e.g., mass, volume).

5. Mass:

• Definition: Mass is a measure of the amount of matter in an object, typically measured in grams (g) or kilograms (kg).

• Representation: Mass is often represented as "m" in formulas.

• Calculation: Mass can be calculated by dividing weight by the acceleration due to gravity (m = W/g). Alternatively, it can be measured using a balance.

6. Volume:

• Definition: Volume is the amount of space an object or substance occupies, typically measured in liters (L), cubic centimeters (cm³), or cubic meters (m³).

• Representation: Volume is represented as "V" in formulas.

• Calculation: Volume can be calculated using the formula for the shape of the object (e.g., V = l × w × h for a rectangular prism).

7. Density:

• Definition: Density is the mass per unit volume of a substance, usually measured in grams per cubic centimeter (g/cm³) or kilograms per liter (kg/L).

• Representation: Density is represented as "ρ" in formulas.

• Calculation: Density is calculated using the formula:Density(ρ)=Mass (m)Volume (V)Density(ρ)=Volume (V)Mass (m)​

8. Using Density to Identify an Unknown Substance:
Density can be used to identify an unknown substance by comparing its calculated density to known values of substances. If the densities match, the substance is identified.

9. Calculate % Error:
The percentage error is a way of comparing the difference between the measured value and the accepted value. It is calculated as:

% Error=(∣Measured Value−Accepted Value∣Accepted Value)×100% Error=(Accepted Value∣Measured Value−Accepted Value∣​)×100

10. Evaluate Instruments to Make the Best Measurements:
To make the best measurements, consider:

• Accuracy: How close the measurement is to the true value.

• Precision: How consistent the measurements are.

• Resolution: The smallest difference the instrument can detect.

11. Convert Between Metric Units:
To convert between metric units, use the metric system’s prefixes (e.g., milli-, centi-, kilo-) and move the decimal point appropriately:

• 1 kilometer (km) = 1,000 meters (m)

• 1 meter (m) = 100 centimeters (cm)

• 1 liter (L) = 1,000 milliliters (mL)

12. Properties of the States of Matter:

• Solid: Particles are closely packed, vibrating in place, with a fixed shape and volume.

• Liquid: Particles are close but can move around, giving it a fixed volume but variable shape.

• Gas: Particles are far apart and move freely, with neither a fixed shape nor volume.

Unit 2

1. Atom and Element:

• Atom: The smallest unit of an element that retains the properties of that element. It consists of a nucleus (protons and neutrons) and electrons orbiting the nucleus.

• Element: A substance made up of only one type of atom, characterized by its number of protons (atomic number). Examples include hydrogen (H), oxygen (O), and carbon (C).

2. Subatomic Particles:

• Protons: Positively charged particles found in the nucleus. The number of protons determines the atomic number.

• Neutrons: Neutral particles found in the nucleus. Neutrons help stabilize the atom and add to its mass.

• Electrons: Negatively charged particles that orbit the nucleus. Electrons are equal in number to protons in a neutral atom.

3. Why Atoms Are Neutral: Atoms are neutral because the number of protons (positive charge) equals the number of electrons (negative charge), canceling out the charges. Therefore, the atom has no overall charge.

4. Isotopes and Their Uses:

• Isotopes: Atoms of the same element with the same number of protons but different numbers of neutrons. This results in different mass numbers.

  • Example: Carbon-12 and Carbon-14 are isotopes of carbon.

• Uses: Isotopes have various applications, such as in medicine (e.g., Carbon-14 for dating fossils), in industry (e.g., Cobalt-60 for cancer treatment), and in research (e.g., tracing reactions in chemistry).

5. Average Atomic Mass:

• Definition: The average mass of an element’s atoms, calculated by considering the relative abundance of each isotope and their respective masses.

• Calculation:Average Atomic Mass=∑(fractional abundance of isotope)×(mass of isotope)Average Atomic Mass=∑(fractional abundance of isotope)×(mass of isotope)

6. Bohr Models for Elements (up to Ca):

• Bohr Model: A representation of an atom with electrons in orbits around the nucleus. Each orbit corresponds to a different energy level.

  • Example for Carbon (C): 6 protons, 6 neutrons, and 6 electrons. Its Bohr model has 2 electrons in the first shell and 4 in the second shell.

7. Electronic Configuration (E.C.) for an Atom:

• Definition: The arrangement of electrons in the orbitals around the nucleus. It follows the Aufbau principle, Pauli exclusion principle, and Hund's rule.

  • Example for Oxygen (O): E.C. = 1s² 2s² 2p⁴.

  • Example for Calcium (Ca): E.C. = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s².

8. Arrangement of the Periodic Table (PT):

• Atomic Number: Elements are arranged in order of increasing atomic number (number of protons).

• Periods: Horizontal rows in the PT. Each period represents a new energy level of electrons.

• Groups: Vertical columns in the PT. Elements in the same group have similar chemical properties due to having the same number of valence electrons.

9. Using the Electronic Configuration (E.C.) to Determine an Element’s Group and Period:

• Group: The group number corresponds to the number of valence electrons (for main group elements).

  • Example: The E.C. of sodium (Na) is 1s² 2s² 2p⁶ 3s¹, so it is in Group 1 (alkali metals).

• Period: The period number corresponds to the highest energy level occupied by electrons.

  • Example: Calcium (Ca) has an E.C. of 1s² 2s² 2p⁶ 3s² 3p⁶ 4s², so it is in Period 4.

10. Using the Location in the PT to Determine Whether an Element Is a Metal or Non-metal:

• Metals: Located on the left side of the PT, they tend to be shiny, malleable, and conductive (e.g., sodium, iron).

• Non-metals: Located on the right side of the PT, they tend to be brittle (solid form), not conductive, and often form negative ions (e.g., oxygen, sulfur).

11. Typical Properties of Metals and Non-metals:

• Metals:

  • Good conductors of heat and electricity.

  • Malleable (can be hammered into sheets).

  • Ductile (can be drawn into wires).

  • Usually have high melting points and are shiny.

• Non-metals:

  • Poor conductors of heat and electricity.

  • Brittle (if solid).

  • Often have low melting and boiling points.

  • Can be gases, liquids, or solids.

12. Elements in Groups 1, 7, 8, and Transition Metals:

• Group 1 (Alkali Metals): Highly reactive metals with one valence electron (e.g., lithium, sodium, potassium).

• Group 7 (Halogens): Highly reactive non-metals with seven valence electrons (e.g., fluorine, chlorine, bromine).

• Group 8 (Noble Gases): Inert gases with eight valence electrons, stable and non-reactive (e.g., helium, neon, argon).

• Transition Metals: Elements in the middle of the PT that typically have high melting points, are good conductors, and can form colorful compounds (e.g., iron, copper, gold).

Unit 3A

1. Ion, Cation, Anion, and Ionic Bond:

• Ion: An atom or molecule that has gained or lost one or more electrons, resulting in a charge.

• Cation: A positively charged ion, formed when an atom loses electrons (e.g., Na⁺, Ca²⁺).

• Anion: A negatively charged ion, formed when an atom gains electrons (e.g., Cl⁻, O²⁻).

• Ionic Bond: The electrostatic force of attraction between oppositely charged ions (cations and anions). This bond forms when one atom donates electrons to another atom, resulting in a stable electron configuration.

2. Types of Ions Formed by Metals and Nonmetals:

• Metals: Metals tend to lose electrons to form cations. For example, sodium (Na) becomes Na⁺ by losing one electron.

• Nonmetals: Nonmetals tend to gain electrons to form anions. For example, chlorine (Cl) becomes Cl⁻ by gaining one electron.

3. Why Ions Are Charged: Ions are charged because the number of electrons does not equal the number of protons. If electrons are lost, the ion becomes positive (cation), and if electrons are gained, the ion becomes negative (anion).

4. Using the EC or PT to Determine the Charge on the Ion Formed:

• EC (Electronic Configuration): The charge can be predicted based on how many electrons an atom needs to gain or lose to achieve a stable electron configuration (often a noble gas configuration).

  • For example, sodium (Na), with an electron configuration of 1s² 2s² 2p⁶ 3s¹, loses one electron to form Na⁺.

  • Chlorine (Cl), with an electron configuration of 1s² 2s² 2p⁶ 3s² 3p⁵, gains one electron to form Cl⁻.

• Periodic Table (PT): The charge on ions can also be predicted from the group number.

  • Elements in Group 1 (alkali metals) form cations with a +1 charge.

  • Elements in Group 17 (halogens) form anions with a -1 charge.

5. Formation of an Ionic Bond:

• Explanation: An ionic bond forms when a metal atom loses electrons (forming a cation) and a non-metal atom gains those electrons (forming an anion). The oppositely charged ions attract each other, creating a stable compound.

  • Example: Sodium (Na) loses one electron to form Na⁺, and chlorine (Cl) gains one electron to form Cl⁻. The resulting Na⁺ and Cl⁻ ions are attracted to each other, forming the ionic compound NaCl (sodium chloride).

• Illustration:Na→Na++e−Na→Na++e−Cl+e−→Cl−Cl+e−→Cl−Na+ and Cl−→NaClNa+ and Cl−→NaCl

6. Deducing the Names and Formulas of Ionic Compounds: To name and write formulas for ionic compounds:

• Name: The cation (metal) name comes first, followed by the anion (non-metal) name with the ending changed to "-ide".

  • Example: Na⁺ and Cl⁻ form NaCl (sodium chloride).

• Formula: The number of each ion is adjusted to ensure the compound is neutral (total positive charge equals total negative charge).

  • Example: Na⁺ and Cl⁻ form NaCl, and Mg²⁺ and Cl⁻ form MgCl₂.

7. Covalent Bond:

• Definition: A covalent bond is a chemical bond formed when two atoms share one or more pairs of electrons. Covalent bonds typically form between nonmetals.

  • Example: In a water molecule (H₂O), the hydrogen and oxygen atoms share electrons.

8. Formation of a Covalent Bond:

• Explanation: In a covalent bond, atoms share electrons to achieve a full outer electron shell (usually 8 electrons for most elements).

  • Example: Two hydrogen atoms (H) each have one electron in their outer shell. By sharing their electrons, they both achieve a stable configuration (2 electrons in the case of hydrogen), forming H₂.

• Illustration:H+H→H2H+H→H2​ (Hydrogen atoms share electrons to form the H₂ molecule).

9. Deducing the Names and Formulas of Covalent Compounds: To name and write formulas for covalent compounds:

• Name: The first element is named first (with a prefix indicating the number of atoms), followed by the second element with an "-ide" ending.

  • Example: CO₂ is carbon dioxide (one carbon, two oxygens).

• Formula: Use prefixes to indicate the number of atoms of each element in the compound.

  • Example: Dinitrogen tetroxide is N₂O₄.

10. Properties of Ionic, Covalent, and Metallic Compounds:

• Ionic Compounds:

  • High melting and boiling points.

  • Conduct electricity when molten or dissolved in water.

  • Soluble in water.

  • Typically form crystalline solids.

• Covalent Compounds:

  • Low melting and boiling points.

  • Do not conduct electricity in any state.

  • Often insoluble in water but soluble in organic solvents.

  • Can exist as gases, liquids, or solids.

• Metallic Compounds:

  • High melting and boiling points.

  • Conduct electricity in solid and molten states.

  • Malleable and ductile (can be hammered into sheets or drawn into wires).

  • Usually shiny and have high tensile strength.

11. How Alloys Change the Properties of Metals:

• Alloy: An alloy is a mixture of two or more elements, usually with metals as the main component.

  • Explanation: Alloys are made to improve or modify the properties of the base metal, such as strength, resistance to corrosion, or hardness.

  • Example: Steel is an alloy of iron and carbon, which is stronger and more durable than pure iron.