Unit 7 - Atomic Structure - Dr. G's...

Atomic Structure & Periodic Table

Unit 7

Ground State vs. Excited State Topic 2

• Chem Catalyst!!

• What do you notice about he speed of the electron as it orbits close to the nucleus compared to further away?

• Which orbit has the most energy? Which has the least energy?

• Is energy lost or gained as the electron moves from an inner to outer energy level?

• Is energy lost or gained as the electron moves from an outer to inner energy level?

• In nature, systems tend toward low energy. Which orbit would the electron want to be in?

Close = slow, far = fast

outside inside

gained (absorbed)

lost (released)

the inside orbit

http://www.mhhe.com/physsci/astronomy/applets/Bohr/applet_files/Bohr.html

Atomic Models

Orbitals are sometimes called electron clouds because they do not have sharp boundaries. Orbitals show where electrons are most likely to be.

Figure 16According to Bohr’s modelof the atom, electrons travel around the nucleus in specific energy levels.

Figure 17According to the currentmodel of the atom, electronsare found in orbitals.

orbital

a region in an atom where thereis a high probability of findingelectrons

Bohr’s Model Confines Electrons to Energy LevelsAccording to Bohr’s model, electrons can be only certain distances fromthe nucleus. Each distance corresponds to a certain quantity of energy thatan electron can have.An electron that is as close to the nucleus as it can beis in its lowest energy level.The farther an electron is from the nucleus, thehigher the energy level that the electron occupies.The difference in energybetween two energy levels is known as a quantum of energy.

The energy levels in Bohr’s model can be compared to the rungs of aladder. A person can go up and down the ladder only by stepping on therungs. When standing on the first rung, the person has the lowest poten-tial energy. By climbing to the second rung, the person increases his or herpotential energy by a fixed, definite quantity. Because the person cannotstand between the rungs on the ladder, the person’s potential energy can-not have a continuous range of values. Instead, the values can be only cer-tain, definite ones. In the same way, Bohr’s model states that an electroncan be in only one energy level or another, not between energy levels.Bohr also concluded that an electron did not give off energy while in agiven energy level.

Electrons Act Like Both Particles and WavesThomson’s experiments demonstrated that electrons act like particlesthat have mass. Although the mass of an electron is extremely small, elec-trons in a cathode ray still have enough mass to turn a paddle wheel.

In 1924, Louis de Broglie pointed out that the behavior of electronsaccording to Bohr’s model was similar to the behavior of waves. For exam-ple, scientists knew that any wave confined in space can have only certainfrequencies. The frequency of a wave is the number of waves that passthrough a given point in one second. De Broglie suggested that electronscould be considered waves confined to the space around a nucleus. Aswaves, electrons could have only certain frequencies. These frequenciescould correspond to the specific energy levels in which electrons are found.

Other experiments also supported the wave nature of electrons. Likelight waves, electrons can change direction through diffraction. Diffractionrefers to the bending of a wave as the wave passes by the edge of an object,such as a crystal. Experiments also showed that electron beams, like waves,can interfere with each other.

Figure 17 shows the present-day model of the atom, which takes intoaccount both the particle and wave properties of electrons. According tothis model, electrons are located in regions around a nucleus thatcorrespond to specific energy levels. Orbitals are regions where electronsare likely to be found. Orbitals are sometimes called electron cloudsbecause they do not have sharp boundaries. When an orbital is drawn, itshows where electrons are most likely to be. Because electrons can be inother places, the orbital has a fuzzy boundary like a cloud.

As an analogy to an electron cloud, imagine the spinning blades of afan. You know that each blade can be found within the spinning imagethat you see. However, you cannot tell exactly where any one blade is ata particular moment.

orbitals,

Atoms and Moles 91

Copyright © by Holt, Rinehart and Winston. All rights reserved.

Figure 16According to Bohr’s modelof the atom, electrons travel around the nucleus in specific energy levels.

Figure 17According to the currentmodel of the atom, electronsare found in orbitals.

orbital

a region in an atom where thereis a high probability of findingelectrons

Bohr’s Model Confines Electrons to Energy LevelsAccording to Bohr’s model, electrons can be only certain distances fromthe nucleus. Each distance corresponds to a certain quantity of energy thatan electron can have.An electron that is as close to the nucleus as it can beis in its lowest energy level.The farther an electron is from the nucleus, thehigher the energy level that the electron occupies.The difference in energybetween two energy levels is known as a quantum of energy.

The energy levels in Bohr’s model can be compared to the rungs of aladder. A person can go up and down the ladder only by stepping on therungs. When standing on the first rung, the person has the lowest poten-tial energy. By climbing to the second rung, the person increases his or herpotential energy by a fixed, definite quantity. Because the person cannotstand between the rungs on the ladder, the person’s potential energy can-not have a continuous range of values. Instead, the values can be only cer-tain, definite ones. In the same way, Bohr’s model states that an electroncan be in only one energy level or another, not between energy levels.Bohr also concluded that an electron did not give off energy while in agiven energy level.

Electrons Act Like Both Particles and WavesThomson’s experiments demonstrated that electrons act like particlesthat have mass. Although the mass of an electron is extremely small, elec-trons in a cathode ray still have enough mass to turn a paddle wheel.

In 1924, Louis de Broglie pointed out that the behavior of electronsaccording to Bohr’s model was similar to the behavior of waves. For exam-ple, scientists knew that any wave confined in space can have only certainfrequencies. The frequency of a wave is the number of waves that passthrough a given point in one second. De Broglie suggested that electronscould be considered waves confined to the space around a nucleus. Aswaves, electrons could have only certain frequencies. These frequenciescould correspond to the specific energy levels in which electrons are found.

Other experiments also supported the wave nature of electrons. Likelight waves, electrons can change direction through diffraction. Diffractionrefers to the bending of a wave as the wave passes by the edge of an object,such as a crystal. Experiments also showed that electron beams, like waves,can interfere with each other.

Figure 17 shows the present-day model of the atom, which takes intoaccount both the particle and wave properties of electrons. According tothis model, electrons are located in regions around a nucleus thatcorrespond to specific energy levels. Orbitals are regions where electronsare likely to be found. Orbitals are sometimes called electron cloudsbecause they do not have sharp boundaries. When an orbital is drawn, itshows where electrons are most likely to be. Because electrons can be inother places, the orbital has a fuzzy boundary like a cloud.

As an analogy to an electron cloud, imagine the spinning blades of afan. You know that each blade can be found within the spinning imagethat you see. However, you cannot tell exactly where any one blade is ata particular moment.

orbitals,

Atoms and Moles 91


Atomic Structure and Light???

Atoms (and specifically the electrons) absorb or emit energy under different conditions when they move to different energy levels.

We can use atomic orbital theory (simplified Bohr’s model) to explain energy changes.

Remember back to our discussion of energy levels of electrons (i.e. valence electrons). Now, let’s expand that discussion to exciting electrons!!

When

writing an excited

state, keep it

simple!

Ground vs. Excited State

Ground State: most probable electron location (lowest energy level possible).

Excited State: place an electron goes after gaining a specific amount of energy.

UNIT 7 TOPIC 2 – GROUND VS. EXCITED STATE

Ground State – most probable electron location (lowest energy level possible)

Excited State – place an electron goes after gaining a specific amount of energy

What happens to the electron configuration of Aluminum if an electron from the

second energy level jumps to the 3rd

energy level?

Ground State –

Excited State –

Which one of these electron configurations would you consider excited?

2-8-1 2-8-1-1 2-7-6 2-5 2-7-9-1

Is energy released when the electron jumps from ground to excited, OR falls

back down from excited to ground? _________________________________

EXPLAIN

Only ONE possible ground state, but MANY possible excited states!

Examples of Excited States!

What happens to the electron configuration of Aluminum if an electron from the second energy level jumps to the 3rd energy level?

Ground State:

!

!

Excited State:

2-8-3

2-7-4

Pick out the Excited State Configurations

2-8-1

2-8-1-1

2-7-6

2-5

2-7-9-1

Is energy released when the electron jumps from the ground state to the excited state, OR when the electron falls back down from the excited state to the ground state? ___________________

Ground

Excited

Excited

Ground

Excited

FALLS BACK DOWN!

Electrons and Light

Light can be thought of as moving waves that have given frequencies, speeds, and wavelengths.

Light travels at 2.998 x 108 m/s. At this speed, light waves take only 500 seconds to travel the 150 million kilometers between the sun and Earth.

The wavelength of the light is the distance between two consecutive peaks or troughs of a wave. The wavelength can vary from 105 m to less than 10-13 m.

This broad range of wavelengths makes up the electromagnetic spectrum.

Electromagnetic Spectrum

Figure 18The electromagnetic spectrum is composed oflight that has a broad rangeof wavelengths. Our eyes candetect only the visible spectrum.

electromagnetic spectrum

all of the frequencies or wavelengths of electromagnetic radiation

1018 Hz 1016 Hz 1014 Hz 1012 Hz 1010 Hz 100 MHz 1 MHz 100 kHz1019 Hz 1017 Hz 1015 Hz 1013 Hz 1011 Hz 109 Hz 10 MHz

0.1 nm

Violet Blue Green Yellow Orange Red

! rays X rays Ultraviolet Infrared Microwaves Radio waves

500 nm 600 nm 700 nm

10 nm 0.1 mm 10 mm 1 m 0.1 km 10 km10 pm 1 nm 0.1 "m 10 "m 1 mm 0.1 m 10 m 1 km

Wavelength

Frequency

Visible spectrum

Electromagnetic spectrum

1 "m

400 nm

1 pm

www.scilinks.orgTopic : Electromagnetic

SpectrumSciLinks code: HW4048

92 Chapter 3

Electrons and LightBy 1900, scientists knew that light could be thought of as moving wavesthat have given frequencies, speeds, and wavelengths.

In empty space, light waves travel at 2.998 × 108 m/s.At this speed, lightwaves take only 500 s to travel the 150 million kilometers between the sunand Earth. The wavelength is the distance between two consecutive peaksor troughs of a wave. The distance of a wavelength is usually measured inmeters. The wavelength of light can vary from 105 m to less than 10−13 m.This broad range of wavelengths makes up the which is shown in Figure 18. Notice in Figure 18 that our eyes are sensitiveto only a small portion of the electromagnetic spectrum. This sensitivityranges from 700 nm, which is about the value of wavelengths of red light, to400 nm, which is about the value of wavelengths of violet light.

In 1905, Albert Einstein proposed that light also has some propertiesof particles. His theory would explain a phenomenon known as thephotoelectric effect. This effect happens when light strikes a metal andelectrons are released. What confused scientists was the observation thatfor a given metal, no electrons were emitted if the light’s frequency wasbelow a certain value, no matter how long the light was on. Yet if lightwere just a wave, then any frequency eventually should supply enoughenergy to remove an electron from the metal.

Einstein proposed that light has the properties of both waves and par-ticles. According to Einstein, light can be described as a stream of particles,the energy of which is determined by the light’s frequency. To remove anelectron, a particle of light has to have at least a minimum energy and there-fore a minimum frequency.

electromagnetic spectrum,


Wavelength vs. Frequency

Frequency and Wavelength are inversely related.

Figure 19The frequency and wavelength of a wave are inversely related.As frequency increases,wavelength decreases.

Red lightLow frequencyLong wavelength

Violet lightHigh frequencyShort wavelength

www.scilinks.orgTopic : Light and ColorSciLinks code: HW4075

www.scilinks.orgTopic : Producing LightSciLinks code: HW4099

Light Is an Electromagnetic WaveWhen passed through a glass prism, sunlight produces the visible spec-trum—all of the colors of light that the human eye can see. You can seefrom Figure 18 on the previous page that the visible spectrum is only a tinyportion of the electromagnetic spectrum. The electromagnetic spectrumalso includes X rays, ultraviolet and infrared light, microwaves, and radiowaves. Each of these electromagnetic waves is referred to as light, althoughwe cannot see these wavelengths.

Figure 19 shows the frequency and wavelength of two regions of thespectrum that we see: red and violet lights. If you compare red and violetlights, you will notice that red light has a low frequency and a long wave-length. But violet light has a high frequency and a short wavelength. Thefrequency and wavelength of a wave are inversely related.

Light EmissionWhen a high-voltage current is passed through a tube of hydrogen gas atlow pressure, lavender-colored light is seen. When this light passesthrough a prism, you can see that the light is made of only a few colors.This spectrum of a few colors is called a line-emission spectrum.Experiments with other gaseous elements show that each element has aline-emission spectrum that is made of a different pattern of colors.

In 1913, Bohr showed that hydrogen’s line-emission spectrum couldbe explained by assuming that the hydrogen atom’s electron can be in anyone of a number of distinct energy levels. The electron can move from alow energy level to a high energy level by absorbing energy. Electrons ata higher energy level are unstable and can move to a lower energy levelby releasing energy. This energy is released as light that has a specificwavelength. Each different move from a particular energy level to a lowerenergy level will release light of a different wavelength.

Bohr developed an equation to calculate all of the possible energies ofthe electron in a hydrogen atom. His values agreed with those calculatedfrom the wavelengths observed in hydrogen’s line-emission spectrum. Infact, his values matched with the experimental values so well that hisatomic model that is described earlier was quickly accepted.

Atoms and Moles 93


Light is a Packet of Particles???In 1905, Albert Einstein proposed that light also has some properties of particles. His theory would explain a phenomenon known as the photoelectric effect. This effect happens when light strikes a metal and electrons are released. What confused scientists was the observation that for a given metal, no electrons were emitted if the light’s frequency was below a certain value, no matter how long the light was on. Yet if light were just a wave, then any frequency eventually should supply enough energy to remove an electron from the metal.

Einstein proposed that light has the properties of both waves and particles. According to Einstein, light can be described as a stream of particles, the energy of which is determined by the light’s frequency. To remove an electron, a particle of light has to have at least a minimum energy and therefore a minimum frequency.

Light Provides Information about Electrons

Normally, if an electron is in a state of lowest possible energy, it is in a ground state. If an electron gains energy, it moves to an excited state.

An electron in an excited state will release a specific quantity of energy as it quickly “falls” back to its ground state.

This energy is emitted as certain wavelengths of light, which give each element a unique line-emission spectrum.

Bright-Line Spectra

The specific set of frequencies of energy emitted by atoms as they return to the ground state produces a unique light, which can be separated into a bright-line spectrum. This can be either an emission spectrum or an absorption spectrum, depending if the gas is hot or cold.

Spectra (the plural of Spectrum)

Bright-Line Spectra

The specific set of frequencies of energy emitted by atoms as they return to the

ground state produces a unique light, which can be separated into a bright-line

spectrum. This can be either an emission spectrum or an absorption spectrum,

depending if the gas is hot or cold.

This is what a real emission spectrum looks like:

Like a prism

Real Example of an Emission Spectrum

Here is the bright-line spectrum for Argon (Ar)

http://www.colorado.edu/physics/2000/quantumzone/

http://www.colorado.edu/physics/2000/quantumzone/

http://www.ifa.hawaii.edu/~barnes/ASTR110L_F03/spectralab_fig2_H.jpg

Other Bright-Line SpectraEvery element has its own unique spectrum, and could be thought of as the “fingerprint” of that element. A spectrum of samples of unknown composition can be observed in order to identify the elements present.

Hydrogen

Helium

Neon

Mercury

Sodium

So what???

Much of what we know about the structure and history of the universe has been figured out by observing the spectra of stars. We know what gases are present in distant stars and galaxies by splitting the light they produce into spectra, and using the elements’ fingerprints to identify them.

Check for Understanding - How all this Works!!

As electrons go from ground state to excited state, they must ____________ (absorb or release) energy.

This process is ____________________ (endo- or exothermic). The opposite is true as electrons “de-excite” or relax back to the ground state. They must give up or ___________ (emit or absorb) energy.

Since this results in the electron (atom) giving up energy, it is _______________________ (endo- or exothermic). The energy is observed as some color of ____________.

If examined through a spectroscope, the element’s characteristic ___________-line spectrum will be observed, and may serve to identify the element.

Absorb

Endothermic

Emit

ExothermicLight

Bright

Regents Practice #1UNIT 7 TOPIC 2 – GROUND VS. EXCITED STATE

REGENTS PRACTICE

An electron in an atom moves from the ground state to an excited state

when the energy of the electron

(1) decreases

(2) increases

(3) remains the same

Base your answers to these questions on the diagram below, which

represents an atom of magnesium-26 in the ground state.

What is the total number of valence electrons in an atom of Mg-26 in the

ground state?

On the diagram below, write an appropriate number of electrons in each

shell to represent a Mg-26 atom in an excited state. Your answer may

include additional shells.

ELABORATE

Regents Practice #2

Base your answers to these questions on the diagram below, which represents an atom of magnesium-26 in the ground state.

!

!

!

What is the total number of valence electrons in an atom of Mg-26 in the ground state?


REGENTS PRACTICE



(1) decreases

(2) increases





ground state?




ELABORATE

2 There are (2-8-2) = 12 TOTAL electrons

Regents Practice #3

On the diagram below, write an appropriate number of electrons in each shell to represent a Mg-26 atom in an excited state. Your answer may include additional shells. (Many possible answers.)


REGENTS PRACTICE



(1) decreases

(2) increases





ground state?




ELABORATE

Weighted Atomic Mass & Allotropes Topic 3

Look at the table of grades below. How could you go about finding this person’s average?

UNIT 7 TOPIC 3 – WEIGHTED ATOMIC MASS & ALLOTROPES

Look at the list of grades below:

93, 94, 95, 95, 95, 95, 95, 95, 95, 94, 95, 95, 95, 95, 95

What do you expect this person’s average to be close to?

Why?

Look at the table of grades below. How could you go about finding this person’s

average?

Percent of Grade Average

Tests 85% 72.9

Quizzes 10% 94.5

Homework 5% 22.8

ENGAGE

Percent of Grade AverageTests 85% 72.9

Quizzes 10% 94.5

Homework 5% 22.8

Weighted Atomic Mass

# Protons: determine which element an atoms is = ATOMIC NUMBER

# electrons: determine the charge on an atom

Protons = electrons: atoms are neutral

Protons > electrons: ION, positively charged (lost electrons)

Protons < electrons: ION, negatively charged (gained electrons)

# neutrons: determines the mass number of an atom (protons + neutrons).


Mass Number = protons + neutrons for an individual atom. It’s a WHOLE number.

Atomic mass = weighted average of mass of all naturally occurring isotopes of an element. This is the DECIMAL on the periodic table.

Isotopes: 2 atoms of the same element having the different number of neutrons.

NOTE: All atoms of an element are isotopes of that element.


LOOK at your periodic tables ... carbon’s atomic mass is 12.011 amu. SO, what is the mass of carbon’s most abundant isotope? (Yes, we did learn this earlier this year!)

Closest to 12, so C-12 is the most abundant

Abundant means LOTS

Most elements have 1, 2, or 3 naturally occurring isotopes. This means that in any sample of the element, these naturally occurring isotopes are all present typically always in the same % ratio.

Isotope Example

C-12: 6 protons, 6 neutrons (90% abundance)



%abundance means that in a sample of carbon (like a lump of coal or a diamond) 90% of the carbon atoms will be carbon-12, 9% will be carbon-13, and 1% will be carbon-14.

Since not all the atoms in a sample of an element have the same mass, we have to calculate an average atomic mass for the element. The average atomic mass is calculated taking into account the different percents of each isotope present.

Average Atomic Mass Example

Here’s the equation (it’s NOT in Table T!!)

!

!

Let’s calculate the atomic mass of carbon!

For example:

The element carbon has three isotopes:

C12

6 = 6 neutrons and is called Carbon -12 isotope 90% abundance

C13

6 = 7 neutrons and is called carbon-13 isotope 9% abundance

C14

6 = 8 neutrons and is called carbon-14 isotope 1% abundance

% abundance means that in a sample of carbon (like a lump of coal or a diamond)

90% of the carbon atoms will be carbon-12, 9% will be carbon-13 and 1% will be

carbon-14.

Since not all the atoms in a sample of an element have the same mass, we have to

calculate an average atomic mass for the element. The average atomic mass is

calculated taking into account the different percents of each isotope present.

Here’s the equation (it’s NOT on Table T!)

Atomic mass = (% isotope 1 x mass isotope 1) + (% isotope 2 x mass Isotope 2) + (% isotope 3 x mass Isotope 3)

100

So, let’s calculate the atomic mass of carbon:

Allotropes

What does this word sound like? Why do you think many people get this confused?

Allotropes: When elements exist in 2 or more different forms. Same atoms, same element, different structure, different formula.

Only a few elements exist like this. Carbon Oxygen Sulfur Phosphorus

Allotrope Examples

Carbon

Allotrope Examples

Oxygen

Allotropes

What does this word sound like? Why do you think many people get this confused?

Allotropes = When elements exist in 2 or more different forms.

Same atoms, same element, different structure, different formula.

Carbon:

Oxygen:

Sulfur & Phosphorus:

Allotrope Examples

Sulfur & Phosphorus

Regents PracticeUNIT 7 TOPIC 3 – WEIGHTED ATOMIC MASS & ALLOTROPES

REGENTS PRACTICE

The weighted atomic mass of a substance is

1. the most abundant isotope

2. The average of all naturally occuring isotopes

3. The atomic number of an element

4. Always the same as the number of protons plus neutrons

ELABORATE

Trends of the Periodic Table Topic 4

• Take out your Periodic Table ... What happens to:

• # of valence electrons as you go left to right across a period?

• # of valence electrons as you go down a group?

• # of energy levels as you go left to right across a period?

• # of energy levels as you go down a group?

• Atomic radius (size) as you go left to right across a period? (Table S)

• Atomic radius (size) as you down a group? (Table S)

• Metallic character as you go left to right across a period?

• Metallic character as you go down a group?

Trends you NEED to KNOW!!

Atomic Radius (Table S)

Electronegativity (Table S)

Ionization Energy (Table S)

Metallic Character

Reactivity (different for metals & nonmetals)

Atomic Radius: the “Falling Snowman”A

tom

ic s

ize

incr

ease

s

Atomic size decreases

It’s ELECTRONEGATIVITY... man!

• Force of attraction the atoms has for an electron. “Desire” for electrons.

• Which do you think would have greater electronegativity; metals or nonmetals? _____________ Why??

• Look at Table S. Which element has the highest electronegativity?

• Why?

Nonmetals

Fluorine only wants 1 e-, so strong desire, and small so valence shell close to pull of electrons

A 3-Dimensional Approach to Electronegativity

Why do noble gases have ZERO electronegativity?

Ionization Energy

• the amount of energy needed to remove an electron from the atom. (The more energy, the harder to remove.)

• Which do you think would have higher ionization energy; metals or nonmetals? ___________ Why??

• Which element has the highest ionization energy? _____

• Why??

Nonmetals

He

Full shell, small size, so e- close to

nucleus & held tightly.

VERY STABLE

A 3-Dimensional Look at Ionization EnergyD

ecre

ases

Increases

Metallic Character

• Describes how readily an atom will ________ electrons to become an ion.

• Let’s refresh our memory about the properties of metals vs. nonmetals...

lose

Metals Nonmetals

•conduct electricity

•shiny, reflective luster

•malleable

•want to lose electrons

•do not conduct

•dull or pearly / no luster

•brittle

•want to gain electrons

Metallic Character Trends

• Based on what you already know about the periodic table, identify the trend in metallic character as you go down a group and from left to right across a period.

Reactivity - Metals

• For METALS, they are more reactive if they ______ electrons easily.

• So, which element is the MOST reactive METAL? _____ Why?

• Group 1: only wants to lose 1 e- to become stable. REALLY wants to lose it.

• Largest group 1, so e- to lose is farthest away from pull of the nucleus.

lose

Fr

Reactivity - Nonmetals

• For NONMETALS, they are more reactive if they ______ electrons easily.

• So, which element is the MOST reactive NONMETAL? ___ Why?

• Halogen, only needs 1 e- to become stable so it REALLY wants one.

• Smallest halogen, so e- to gain is closest to the pull of the positively charged nucleus.

gain

F

Reactivity TrendsIn

crea

ses

met

al re

activ

ity

Metal reactivity increases

Nonmetal reactivity increases

Non

met

al re

activ

ity

incr

ease

s

Regents PracticeUNIT 7 TOPIC 4 – TRENDS OF THE PERIODIC TABLE

REGENTS PRACTICE

Atoms of which element have the greatest tendency to gain electrons?

A) Bromine

B) Fluorine

C) Chlorine

D) Iodine

Which element in Group 15 has the strongest metallic character?

A) As

B) P

C) N

D) Bi

Compared to atoms of metals, atoms of nonmetals generally

A) Have lower first ionization energies

B) Conduct electricity more readily

C) Have higher electronegativities

D) Lose electrons more readily

ELABORATE

UNIT 7 TOPIC 4 – TRENDS OF THE PERIODIC TABLE

REGENTS PRACTICE


A) Bromine

B) Fluorine

C) Chlorine

D) Iodine


A) As

B) P

C) N

D) Bi






ELABORATE

UNIT 7 TOPIC 4 – TRENDS OF THE PERIODIC TABLE

REGENTS PRACTICE


A) Bromine

B) Fluorine

C) Chlorine

D) Iodine


A) As

B) P

C) N

D) Bi






ELABORATE

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