Chapter 9 Chem Quiz

Bohr model and the quantum-mechanical model

propose explanations for the inertness of helium, the reactivity of hydrogen, and the periodic law

• They explain how electrons exist in atoms and how those electrons affect the chemical and physical properties of elements

Niels Bohr and Erwin Schrödinger (along with albert einstein)

played a role in the development of quantum mechanics, yet bewildered by their own theory of wave-particle duality for the electron

Light

a form of electromagnetic radiation that travels through space at constant speed of 3.0 × 10⁸ m/s (c) (186,000 mi/s)

Modern Atomic Structure

c= λv E= hv E = hc/λ

c= 3.0 × 10⁸ m/s 

h= 6.6262x 10^-34 Js

Wavelength

λ (lambda); the distance between adjacent wave crests 

Spectrum of color

• White light contains spectrum of wavelength → spectrum of color

• ROYGBIV–seen in a rainbow or when light passes through a prism

  • Red, orange, yellow, green, blue, indigo, violet

• Red light–longest wavelength

• Violet light–shortest wavelength

• Color seen when it is reflected and all other colors are absorbed

Frequency

v nu; the number of cycles or crests that pass through a stationary point in one second

Relationship between wavelength and frequency

inverse proportion (shorter the wavelength, the higher the frequency and vice versa)

Photon

particle of light; single packet of light energy

• Amount of energy depends on wavelength of light–the shorter the wavelength, the greater the energy

• Light waves have more energy when their crest are closer together–higher frequency and shorter wavelength

The Electromagnetic Spectrum

Gamma rays

shortest wavelength and more energetic

• Produced by the sun, stars, and unstable atomic nuclei on Earth

• Excessive human exposure can be dangerous because the high energy of the photons can damage biological molecules

X-rays

pass through many substances that block visible light and are used to image internal bones and organs

• Carry enough energy to damage biological molecules

• Several yearly exposures are pretty harmless, excessive exposure to x-rays increase cancer risk

Ultraviolet or UV light

component of sunlight that produces sunburn or suntan

• Though not as strong as gamma or x rays, has enough energy to damage biological molecules

• Increase risk to skin cancer, cataracts, and causes premature wrinkles

Visible light

ranges from violet to red

• Do NOT damage biological molecules

• Cause molecules in our eyes to rearrange, which sends a signal to our brains that results in vision

Infrared light

felt when hand placed hear a hot object

• Warm abjects (even humans) emit infrared light

• Invisible to our eyes, sensors for it can detect it and used for night vision technology

• Warm objects like humans glow as much as a lightbulb in the visible region of the spectrum

Microwaves

used for radar and in microwave ovens

• Efficiently absorbed by water and can heat substances that contain water

• Substances that contain water (food) are warmed by radiation of a microwave, but substances that don't have water, like a plate, cannot

Radio waves

longest wavelength

• Used to transmit signal used by AM and FM radio, cellular telephones, TV, and other forms of communication

Emission Spectra

White light spectrum is continuous with radiation emitted at every wavelength; the emission spectrum of an individual element includes only certain specific wavelengths

Bohr model

can explain the emission spectrum of hydrogen; each orbit specified by a quantum number (n) which also specifies the orbits energy; cannot predict spectra for atoms with more than one electron

Energy of each Bohr orbit

specified by quantum number n= 1,2,3 is fixed (quantized); like steps of a ladder, each specific distance from nucleus and each at a specific energy

• Impossible for an electron to be between orbits in Bohr model

Excitation and Emission

when a hydrogen atom absorbs a quantum of energy, an electron is excited to a higher energy orbit, then the electron relaxes back down to a lower energy orbit emitting a photon of light

• Since amt of energy in a photon is directly related to its wavelength, the photon has a specific wavelength

• Light emitted by excited atoms consists of specific lines at specific wavelengths, each corresponding to a specific transition between two orbits

  • Ex: the line at 486 nm in the hydrogen emission spectrum corresponds to an electron relaxing from the n = 4 orbit to the n = 2 orbit

  • Ex: line at 657 nm (longer wavelength and lower energy) corresponds to an electron relaxing from the n = 3 orbit to the n = 2 orbit

How was the Bohr model successful? What did it fail to do?

It was successful because it predicted the lines of the hydrogen emission spectrum.

It failed to predict the emission spectra of other elements that contained more than one electron.

Quantum mechanical or wave mechanical model

describes electron orbitals, which are electron probability maps that show the relative probability of finding an electron in various places surrounding the atomic nucleus

• replaced the Bohr model in the early twentieth century

• quantum-mechanical orbitals replaces Bohr obits. 

• can predict the bright-line spectra of other elements

electron configuration

indicates which orbitals are occupied for a particular atom; shows the occupation of orbitals by electrons for a particular atom

Orbitals

represent probability maps that show a statistical distribution of where the electron is likely to be found

• electrons do not behave like particles flying through space

• does NOT represent the exact path that an electron takes as it travels through space

Principle quantum number (n)

specify an orbital (or orbitals) and the principal shell of the orbital; lowest-energy orbital in the quantum-mechanical model is called the 1s orbital

• The higher the principal quantum number, the higher the energy of the orbital

• Possible numbers are n=1, 2, 3 with energy increasing as n increases

Ground state

lowest energy state

Excited state

when an electron is in the higher energy orbital after the absorption of energy by a hydrogen atom that causes the electron to jump (or make a transition) to a higher energy orbital. 

• All atoms have one ground state and many excited states

Subshell

indicated by the letter, specifies shape; s, p, d, f

• s—hold up to 2 electrons in total (1 pair)

• p—hold up to 6 electrons in total (3 pairs)

• d—hold up to 10 electrons total (5 pairs)

• f—hold up to 14 electrons total (7 pairs)

Dot density

proportional to the probability of finding the electron; for the 1s orbital is greatest near the nucleus and decreases farther away from the nucleus–most likely electron found close to the nucleus than far away from it

The Number of Subshells in a Given Principal Shell Is Equal to the Value of n

Electron Configurations in the Periodic Table

1s and 2s orbital shape

2p orbital

3d orbitals

Orbital diagram

shows the electrons as arrows in a box representing the orbital

• Arrow represents electron spin

Pauli exclusion principle

orbitals may hold no more than two electrons with opposing spins; When two electrons occupy the same orbital, they must have opposing spins

Energy of ordering of orbitals for multielectron atoms

Aufbau Principle

1. Electrons enter orbitals of lower energy first

2. The various orbitals within a sublevel have equal energy

3. The “s” sublevel is always the lowest energy

Hund’s rule

When filling orbitals of equal energy, electrons fill them singly first, with parallel spins

Noble Gas Core Notation

When writing electron configurations for elements beyond neon—or beyond any other noble gas—the electron configuration of the previous noble gas can be abbreviated by the symbol for the noble gas in brackets

• Ex: Na: 1s22s22p63s1 → Na: [Ne]3s1

Valence Electrons

are the electrons in the outermost principal shell (the principal shell with the highest principal quantum number, n)

• Involved in chemical bonding

Core electrons

Electrons that are not in the outermost principal shell

Periodic trends in Electron Configurations (main group elements)

• Number of valence electrons is equal to the group number of its column (besides Helium)

• Row number is equal to the number of the highest principal shell

Periodic trends in Electron Configurations (transition metals)

•  The principal quantum number of the d orbital is equal to the row number minus 1

• For the first transition series the outer configuration is 4s²3d^x  (x = number of d electrons)

  • Two exceptions: Cr is 4s¹3d⁵ and Cu is 4s¹3d¹⁰

• The number of outer shell electrons in a transition series does not change as you move across a period

  • transition series represents the filling of core orbitals and the number of outer shell electrons is mostly constant—either 2 or 1. 

  • (2e–) for  4s²3d^x

  • (1e–) for  4s¹3d⁵ or 4s¹3d¹⁰

Write the Electron Configuration for Any Element Based on Its Position in the Periodic Table

• The inner electron configuration is the electron configuration of the noble gas that immediately precedes that element in the periodic table. Represent the inner configuration with the symbol for the noble gas in brackets.

• The outer electrons can be determined from the element’s position within a particular block (s, p, d, or f) in the periodic table. Trace the elements between the preceding noble gas and the element of interest, and assign electrons to the appropriate orbitals.

• The highest principal quantum number (highest n value) is equal to the row number of the element in the periodic table.

• For any element containing d electrons, the principal quantum number (n value) of the outermost d electrons is equal to the row number of the element minus 1

Noble Gases

• Atoms with 8 electrons (2 for helium) are predicted to be low in energy and stable

• Chemically stable, relatively inert (non reactive)

• Electron configurations close to noble gases are most reactive because they can attain noble gas electron configurations by losing or gaining a small number of electrons

Alkali Metals

• Most receive since their outer electron configuration is (ns¹) is 1 electron beyond noble gas configuration

• If react to lose the electron, they attain a noble gas configuration

Alkaline Earth metals

• All have electron configurations ns² and are 2 electrons beyond a noble gas configuration

• Tend to lose 2 electrons, forming 2+ ions and attaining a noble gas configuration

Halogens

• All have ns²np⁵ electron configurations therefore 1 short of a noble gas configuration

• Tend to gain 1 electron, forming 1- ions and attaining a noble gas configuration

atomic size

distance from the nucleus to the outermost electron

Periodic trend: Atomic Size

• As you move to the right across a period in the periodic table, atomic size decreases

  • atomic size of an atom is determined by the distance between the outermost electrons and the nucleus

  • Size of an orbital depends on principle quantum number

  • Each step across period, number of protons increase → greater pull on electrons from the nucleus causing size to decrease

• As you move down a column in the periodic table, atomic size increases

  • the highest principal quantum number, n, increases

  • Since size of an orbital increases with increasing principal quantum number, the electrons that occupy the outermost orbitals are farther from the nucleus a you move down a column

Ionization energy

the amount of energy needed to remove an electron from a neutral atom, making it a positively charged ion

Ionization energy Periodic Trend

increases as you move right across a period, and decreases as you move down a column in the periodic table

Metallic Character

ease with which an atom becomes a cation

Metallic Character Periodic Trend

decreases as you move right across period and increases as you move down a column in the periodic table

• Metals tend to lose electrons in their chemical reactions, while nonmetals tend to gain electrons.

• As you move across a period in the periodic table, ionization energy increases, which means that electrons are less likely to be lost in chemical reactions.

• Metallic character decreases as you move to the right across a period and increases as you move down a column in the periodic table.