Chapter 4, 5, 6 Chem Test Flashcards
conservation of mass
mass is neither created nor destroyed in chemical reactions
ex. Burning paper
Definite proportions
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
Multiple proportions
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
Dalton’s atomic theory
Elements are composed of atoms
Atoms of the same element are identical
Cathode Ray Tube (CRT) Experiment
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.
Gold Foil Experiment (Rutherford Experiment)
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.
Thomson’s model (Plum Pudding Model)
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.
Rutherford’s Model (Nuclear Model)
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.
Bohr’s Model (Bohr Atomic Model)
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.
Electron
Charge: Negative (-1 elementary charge).
Location: Orbits the nucleus in electron shells.
Mass: Very light, approximately 1/1836 times the mass of a proton.
Proton
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.
Neutron
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.
Isotopes
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.
Atomic Number
Atomic number is the number of protons in an atom’s nucleus. It defines the element and determines its chemical properties.
Mass Number
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.
Average Atomic Mass
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.
Cations (Ion)
Cations are positively charged ions formed when an atom loses one or more electrons. They have fewer electrons than protons.
Anions (Ion)
Anions are negatively charged ions formed when an atom gains one or more electrons. They have more electrons than protons.
Calculate the average atomic mass from percent isotopes
convert % to decimal → multiply percent by mass → sum the numbers
Proton (p)
The number of protons in the nucleus is given by the atomic number, which is the lower number in the symbol.
Electron (e)
In a neutral atom, the number of electrons is equal to the number of protons, which can be determined from the atomic number.
Neutron (n)
To find the number of neutrons, subtract the number of protons (atomic number) from the mass number (the upper number in the symbol).
Mass Number (A)
The mass number is the sum of protons and neutrons. It’s the upper number in the symbol.
Charge (Q)
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.
Isotope Status
Isotopes have the same number of protons (same atomic number) but different numbers of neutrons (different mass numbers).
major waves of the electromagnetic spectrum
Radio Waves
Microwaves
Infrared Waves
Visible Light
Ultraviolet (UV) Rays
X-Rays
Gamma Rays
relationship between wavelength and frequency
indirect
energy and frequency changes for Balmer Series
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.
energy and frequency changes for Lyman Series
Energy and Frequency: Transitioning from higher energy levels to the first energy level.
Wavelength: In the ultraviolet range.
energy and frequency changes for Paschen Series
Energy and Frequency: Transitioning from higher energy levels to the third energy level.
Wavelength: In the infrared range.
What is the result when electrons transition from an excited state to a ground state
they release energy in the form of a photon (light), resulting in the emission of light
properties of the four quantum numbers
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).
electron configuration
orbital notations
core (noble gas) notations
1s^2 2s^2 . . .
the squares with up down arrows
[He] . . .
Aufbau Principle
Electrons fill the lowest energy orbitals first before moving to higher energy orbitals as depicted in the Aufbau diagram.
Hund’s Rule
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.
Pauli Exclusion Principle
Each orbital can hold a maximum of two electrons with opposite spins (one “up” and one “down”), as shown in orbital diagrams.
Energy Shells
Energy levels that surround the nucleus and contain subshells. They are represented by the principal quantum number (n).
Subshells (Sublevels)
Subdivisions within energy shells, each with a unique shape (s, p, d, f) represented by the angular momentum quantum number (l).
Orbitals
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)
Electron Numbers
The count of electrons in an atom, which determines how orbitals are filled within subshells and subshells within energy shells
Döbereiner
Introduced the concept of triads, grouping elements with similar properties in sets of three.
Meyer
Recognized the periodicity of elements’ properties and organized them by increasing atomic volume
Newlands
Proposed the Law of Octaves, grouping elements in sets of eight with similar properties.
Mendeleev
Developed the periodic table based on increasing atomic mass and accurately predicted the existence and properties of undiscovered elements.
Ramsay
Discovered noble gases, which were added to the periodic table as a new group.
Moseley
Established the modern periodic table by arranging elements by increasing atomic number, leading to the periodic law.
Seaborg
Extended the periodic table to accommodate the actinides and lanthanides, contributing to the modern periodic table’s completeness.
Down a Group (Column)
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.
Across a Period (Row):
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.
Significance of Isoelectronic Species:
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
Interpreting Size from Isoelectronic Species
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.
Valence Electrons
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.
Valence Numbers
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
