Chapter 4, 5, 6 Chem Test

Question 1

conservation of mass

Answer 1

mass is neither created nor destroyed in chemical reactions
ex. Burning paper

Question 2

Definite proportions

Answer 2

In a chemical compound, elements are always present in fixed mass ratios
ex. Water (H2O) always has a 2:16 mass ratio of hydrogen to oxygen

Question 3

Multiple proportions

Answer 3

When two elements form multiple compounds, their masses combine in simple whole number ratios
ex. Carbon and oxygen form carbon monoxide (CO) and carbon dioxide (CO2) with ratios of 1:1 and 2:1, showing this principle

Question 4

Dalton’s atomic theory

Answer 4

Elements are composed of atoms
Atoms of the same element are identical

Question 5

Cathode Ray Tube (CRT) Experiment

Answer 5

Experiment: J.J. Thomson observed the behavior of cathode rays in a vacuum tube.
Conclusions: Identified electrons as negatively charged subatomic particles with a high charge-to-mass ratio, leading to the “plum pudding” model of the atom.

Question 6

Gold Foil Experiment (Rutherford Experiment)

Answer 6

Experiment: Ernest Rutherford fired alpha particles at a gold foil.
Conclusions: Discovered that atoms are mostly empty space with a dense, positively charged nucleus, leading to the nuclear model of the atom with electrons orbiting the nucleus.

Question 7

Thomson’s model (Plum Pudding Model)

Answer 7

depicted the atom as a positively charged “pudding” with negatively charged electrons (“plums”) embedded within it.
Key Idea: Atoms contain negatively charged electrons distributed within a positively charged matrix.

Question 8

Rutherford’s Model (Nuclear Model)

Answer 8

Description: Ernest Rutherford’s model, proposed in 1911, introduced the concept of a tiny, dense, positively charged nucleus at the center of the atom, with electrons orbiting around it.
Key Idea: The nucleus contains most of the atom’s mass, and electrons orbit around it.

Question 9

Bohr’s Model (Bohr Atomic Model)

Answer 9

Description: Niels Bohr’s model, proposed in 1913, extended Rutherford’s model by introducing quantized energy levels or electron shells where electrons orbit the nucleus.
Key Idea: Electrons orbit the nucleus in specific energy levels, and they can jump between these levels by absorbing or emitting energy.

Question 10

Electron

Answer 10

Charge: Negative (-1 elementary charge).
Location: Orbits the nucleus in electron shells.
Mass: Very light, approximately 1/1836 times the mass of a proton.

Question 11

Proton

Answer 11

Charge: Positive (+1 elementary charge).
Location: Located in the nucleus at the center of the atom.
Mass: Relatively heavy, approximately 1836 times the mass of an electron.

Question 12

Neutron

Answer 12

Charge: Neutral (no net charge).
Location: Also located in the nucleus.
Mass: Similar to that of a proton, approximately 1839 times the mass of an electron.

Question 13

Isotopes

Answer 13

Isotopes are atoms of the same element with the same number of protons (same atomic number) but different numbers of neutrons. They have similar chemical properties but different atomic masses.

Question 14

Atomic Number

Answer 14

Atomic number is the number of protons in an atom’s nucleus. It defines the element and determines its chemical properties.

Question 15

Mass Number

Answer 15

Mass number is the sum of protons and neutrons in an atom’s nucleus. It provides the atom’s mass, but it doesn’t specify the element.

Question 16

Average Atomic Mass

Answer 16

This is the weighted average of the masses of all naturally occurring isotopes of an element. It considers the abundance of each isotope in a sample.

Question 17

Cations (Ion)

Answer 17

Cations are positively charged ions formed when an atom loses one or more electrons. They have fewer electrons than protons.

Question 18

Anions (Ion)

Answer 18

Anions are negatively charged ions formed when an atom gains one or more electrons. They have more electrons than protons.

Question 19

Calculate the average atomic mass from percent isotopes

Answer 19

convert % to decimal → multiply percent by mass → sum the numbers

Question 20

Proton (p)

Answer 20

The number of protons in the nucleus is given by the atomic number, which is the lower number in the symbol.

Question 21

Electron (e)

Answer 21

In a neutral atom, the number of electrons is equal to the number of protons, which can be determined from the atomic number.

Question 22

Neutron (n)

Answer 22

To find the number of neutrons, subtract the number of protons (atomic number) from the mass number (the upper number in the symbol).

Question 23

Mass Number (A)

Answer 23

The mass number is the sum of protons and neutrons. It’s the upper number in the symbol.

Question 24

Charge (Q)

Answer 24

For a neutral atom, the total charge is zero. If the atom has gained or lost electrons, it will carry a charge (±) indicated as superscript.

Question 25

Isotope Status

Answer 25

Isotopes have the same number of protons (same atomic number) but different numbers of neutrons (different mass numbers).

Question 26

major waves of the electromagnetic spectrum

Answer 26

Radio Waves
Microwaves
Infrared Waves
Visible Light
Ultraviolet (UV) Rays
X-Rays
Gamma Rays

Question 27

relationship between wavelength and frequency

Answer 27

indirect

Question 28

energy and frequency changes for Balmer Series

Answer 28

Energy and Frequency: Transitioning from higher energy levels to the second energy level.
Wavelength: In the visible range, mainly in the red and blue regions.

Question 29

energy and frequency changes for Lyman Series

Answer 29

Energy and Frequency: Transitioning from higher energy levels to the first energy level.
Wavelength: In the ultraviolet range.

Question 30

energy and frequency changes for Paschen Series

Answer 30

Energy and Frequency: Transitioning from higher energy levels to the third energy level.
Wavelength: In the infrared range.

Question 31

What is the result when electrons transition from an excited state to a ground state

Answer 31

they release energy in the form of a photon (light), resulting in the emission of light

Question 32

properties of the four quantum numbers

Answer 32

Principal Quantum Number (n): Energy level or shell of an electron.
Angular Momentum Quantum Number (l): Orbital shape or subshell.
Magnetic Quantum Number (ml): Orbital orientation in space.
Spin Quantum Number (ms): Electron’s spin, either up (+1/2) or down (-1/2).

Question 33

electron configuration
orbital notations
core (noble gas) notations

Answer 33

1s^2 2s^2 . . .
the squares with up down arrows
[He] . . .

Question 34

Aufbau Principle

Answer 34

Electrons fill the lowest energy orbitals first before moving to higher energy orbitals as depicted in the Aufbau diagram.

Question 35

Hund’s Rule

Answer 35

When filling degenerate orbitals (same energy level), electrons enter them singly before pairing up. This minimizes electron-electron repulsion and is illustrated in orbital diagrams.

Question 36

Pauli Exclusion Principle

Answer 36

Each orbital can hold a maximum of two electrons with opposite spins (one “up” and one “down”), as shown in orbital diagrams.

Question 37

Energy Shells

Answer 37

Energy levels that surround the nucleus and contain subshells. They are represented by the principal quantum number (n).

Question 38

Subshells (Sublevels)

Answer 38

Subdivisions within energy shells, each with a unique shape (s, p, d, f) represented by the angular momentum quantum number (l).

Question 39

Orbitals

Answer 39

Specific regions within subshells where electrons are likely to be found. Each orbital can hold up to 2 electrons and is defined by the magnetic quantum number (ml)

Question 40

Electron Numbers

Answer 40

The count of electrons in an atom, which determines how orbitals are filled within subshells and subshells within energy shells

Question 41

Döbereiner

Answer 41

Introduced the concept of triads, grouping elements with similar properties in sets of three.

Question 42

Meyer

Answer 42

Recognized the periodicity of elements’ properties and organized them by increasing atomic volume

Question 43

Newlands

Answer 43

Proposed the Law of Octaves, grouping elements in sets of eight with similar properties.

Question 44

Mendeleev

Answer 44

Developed the periodic table based on increasing atomic mass and accurately predicted the existence and properties of undiscovered elements.

Question 45

Ramsay

Answer 45

Discovered noble gases, which were added to the periodic table as a new group.

Question 46

Moseley

Answer 46

Established the modern periodic table by arranging elements by increasing atomic number, leading to the periodic law.

Question 47

Seaborg

Answer 47

Extended the periodic table to accommodate the actinides and lanthanides, contributing to the modern periodic table’s completeness.

Question 48

Down a Group (Column)

Answer 48

Atomic Size: Increases. More energy levels are added, making the atom larger.
Ionization Energy: Decreases. Electrons are farther from the nucleus, so they are easier to remove.
Electron Affinity: Decreases. Atoms are less likely to gain electrons.
Electronegativity: Decreases. Atoms are less likely to attract electrons.
Ionic Radii: Increase for positive ions (cations) and decrease for negative ions (anions).
Metallic Character: Increases. Atoms become more metallic and have a stronger tendency to lose electrons.

Question 49

Across a Period (Row):

Answer 49

Atomic Size: Decreases. More protons in the nucleus pull electrons closer.
Ionization Energy: Increases. It becomes harder to remove electrons.
Electron Affinity: Becomes more negative (tendency to gain electrons).
Electronegativity: Increases. Atoms have a stronger pull on electrons.
Ionic Radii: Decrease for both cations and anions.
Metallic Character: Decreases. Atoms become less metallic and have a stronger tendency to gain electrons.

Question 50

Significance of Isoelectronic Species:

Answer 50

Isoelectronic species are atoms or ions with the same number of electrons. They help us compare how different elements or ions from various elements interact in chemical reactions, particularly in terms of charge and reactivity

Question 51

Interpreting Size from Isoelectronic Species

Answer 51

In isoelectronic species, the size (atomic or ionic) decreases with increasing nuclear charge (the number of protons in the nucleus). Smaller species have a greater nuclear charge, attracting electrons more strongly, resulting in a smaller size.

Question 52

Valence Electrons

Answer 52

Valence electrons are the electrons in the outermost energy level of an atom. They determine the chemical properties and reactivity of an element

They are involved in chemical bonding and reactions. The number of valence electrons influences an element’s ability to form bonds and its position in the periodic table.

Question 53

Valence Numbers

Answer 53

Valence numbers (oxidation numbers) represent the charge an atom assumes in a compound. They are crucial in understanding how elements combine to form compounds and in balancing chemical equations