What happens to the number of energy levels as you go from left to right on the periodic table?

As you move from left to right, the nucleus gains protons.

This increases the positive charge of the nucleus and its attractive force on the electrons.

At the same time, electrons are added to the atoms as you move from left to right across a period.

These electrons reside in the same energy shell and do not offer complete shielding.

The effect of the increasing proton number is greater than that of the increasing electron number.

As a result, the valence electrons are held closer to the nucleus, and the atomic radius decreases.

The chart below shows the relative sizes of atoms.

The following video shows how the atomic radius of elements changes as you move across the Periodic Table.

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We use the periodic table to help us recognize certain trends of physical and chemical properties of the elements. You need to memorize the trends. A trend is generally "it gets bigger" or "it gets smaller" sort of thing. All our trends describe the trend in two directions on the periodic table: 1) across a row, and 2) up and down a column. Here are the important ones for us.

Atomic Size

The smallest atom on the periodic table is helium, He, and has a radius of 31 pm. Yeah, He is even smaller than hydrogen, H, which is 53 pm. Which atom is the largest? That would be cesium, Cs, which comes in with a radius of 343 pm. So that is roughly a 10:1 ratio of largest to smallest. Sometimes we just do a generalized bit of rounding as well and say things like atoms range from about 50 pm to 300 pm which is more of a 6:1 ratio. Oh well, you should just wrap your head around the general range of all atomic sizes the extremes are 31 pm and 343 pm... so chopping that to 50-300 pm isn't a big deal.

Atoms get bigger as you go down a column on the periodic table. This is because in going down a column you are jumping up to the next higher main energy level (n) and each energy level is further out from the nucleus - that is, a bigger atomic radius.

Atoms get smaller as you go across a row from left to right. This may seem counterintuitive but it is the fact. The logic is that as you go across rows, you are staying in the same main energy level (n) so electrons are entering the atomic atmosphere at about the same distance. However, as you go across, the nuclei are getting more and more positive (more protons) - therefore there is more + to – attraction and the electron cloud is pulled in tighter and therefore a smaller radius.

So on any one row, the group 1 atoms (alkali metals) are the biggest on that row and the group 18 atoms (noble gases) are the smallest. Below is a simple graphic illustrating the atomic radii trends.

Monatomic Ion Size

Now that you have the trend for neutral atoms, let's modify or tweak those sizes for when the atom is changed into a cation or anion.

Cations: Metals tend to lose their electrons to make stable cations. The typical number is one to three electrons to make +1, +2, and +3 cations. Realize that when you make a cation from a monatomic neutral species, you are removing electrons from the outmost valence shell. Upon each e– removal, there are fewer e– repulsions which means the remaining electrons are pulled in tighter than before. This means that cations have smaller radii than the neutral atom from which they came from. And, each subsequent removal of additional electrons leads to smaller and smaller cation species. This is illustrated below starting on the left with a neutral atom.

Anions: Non-metals tend to gain electrons to make stable anions. So in a likewise but opposite manner - we ADD electrons to the valence shell thus increasing electron repulsions which means the resulting anion is bigger than the atom from which they came. The more electrons you add, the bigger the anion gets. This is illustrated in the diagram below starting on the left with a neutral atom.

Here's a figure from Wikipedia showing the neutral atomic radii vs the ionic radii sizes for some cations and anions.

Ionization Energy (IE)

Ionization energy is the amount of energy it takes to remove one electron from a neutral atom (A) in order to form a +1 cation. The reaction (with energy shown) is

A   +   energy →   A+   +   e–

The energy needed to do this must overcome the attraction of the outermost electron to the nucleus. All atoms have a wide variety of energies needed to do this, but they DO follow a trend that is easily seen on the periodic table. Much like all the trends, the two extremes of this property are at the bottom left (smallest IE) and the top right (largest IE). Going down a column, IE's decrease. Going across rows, IE's increase.

Electron Affinity (EA)

Electron affinity is the amount of energy released when one electron is added to a neutral atom (A) in order to form a –1 anion. The reaction (with energy shown) is

A   +   e–   →   A–   +   energy

You can think of EA as the "desire of an electron" by an atom. If the atom "wants" the electron a lot, then the EA is big. Less desire is smaller energy and there is even no desire and the numbers go to zero and even negative. The trends on the periodic table are not as pronounced as with other trends (they're a bit janky) - but in general, the upper right corner has the largest EAs while the lower left corner has the lowest values. I'm including this for the purpose of pointing out this is a real measurement and the recognition of EA is more important for our studies than the actual values. Move on to electronegativity now.

Electronegativity (EN)

Electronegativity is a relative scale from zero to four that measures the "desire" or "pull" on electron pairs. Electronegativity is the purposeful human friendly scale from 0 to 4 that electron affinity lacked. The maximum of 4.0 on this scale belongs to fluorine (top right). The minimum of 0.8 on this scale belongs to cesium (bottom left). Think of EN as the "pull" on electron pairs in a molecule by an atom. We use it the most of the three trends/properties last listed. And yes, we ignore the noble gases for EN values because they are happy as is - they have no desire for any shared electrons and they don't form bonds, so no values for them.

We will rarely need the actual numbers for electronegativity. Just knowing approximately which elements are the most electronegative (upper right corner) helps us in recognizing and assigning polarity of bonds and ultimately compounds. The non-metals tend to be at or above 2.0 on the scale which means they "want" electrons far more than all the metals which tend to all be less than 2.0 on the scale. Go to Wikipedia or other online resources if you want the actual numbers for electronegativity.

Summary of ALL Trends

Below is an illustration showing how the extremes of all properties (trends) are in the same two regions.

I will always "ignore" the last row of elements when talking trends. Why?

Many students ask me, "Why did you say cesium is the largest atom instead of francium?". Well, pretty much that entire 7th row of elements are very radioactive. Francium's most stable isotope has a half-life of only 22 minutes. So it doesn't really stick around long enough to really even do any real chemistry. So when I talk about trends, the "extreme" in the bottom left corner is cesium.

Just remember this: when you hear about "general chemistry" or "principles of chemistry", there is a hidden prepositional phrase at the end of that. That phrase is "of the stable elements". So yes, 99% of the time when discussing chemistry of the elements and their trends, only the non-radioactive/stable elements are relevant. IF you study nuclear chemistry in a class, then yes, all those unstable atoms are relevant again. WE are not doing nuclear chemistry in this class or book.

One more thing in the upper right corner... the trends of electron affinity and electronegativity are only relevant to elements that actually react in chemical reactions. So that is why fluorine (not helium or neon) wins the "extreme" trend in the upper right corner of the periodic table with those properties. FYI - helium does win in the ionization energy contest (and smallest atom) because that is the energy to remove an electron - helium is definitely the toughest element to remove an electron from. The other noble gases are very stingy as well.

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When two non-metal elements get together we have a dilemma in a way. BOTH elements by their nature "want" more electrons to achieve noble gas electron configuration - they want 8 electrons in their outer shell as we say which is s2p6. So neither element is going to give up electrons, they are holding on to what they got (thank you Bon Jovi). Or maybe even they're saying to their electrons "Never Gonna Give You Up" as they Rickroll into a covalent bond. The POINT being... they have no problem in sharing electrons so that they can each make it to 8. An example is carbon tetrachloride (CCl4 - we'll get to naming in a bit). The carbon has 4 electrons in its outer shell (2s22p2) and each of the four chlorines has 7 (2s22p5). So what happens is that each chlorine shares one electron with the carbon and the carbon shares one electron with each of the chlorines (that's 4 total). The end result is that there are four covalent bonds between the carbon and the chlorines. All five atoms are holding on tightly to the shared electrons - this is the basis of all covalent bonds. This is shown below with electron dot formulas. Each dot is a valence electron around that atom.

Notice how the finished product has 8 electrons (dots) around each element. This is the end result of elements obeying the Octet Rule. The diagram below emphasizes this fact by circling the 8 electrons around each atom. The other diagram/drawing shows how we swap out a shared pair of dots for a line. The line represents 2 shared electrons and we call it a covalent bond.

This is the way we diagram a molecule with covalent bonds. When you start learning this skill, you will definitely use all the dots and move things around until you find the structure with all 8's surrounding the atoms. We will dive into electron line/dot formulas a little later, for now lets learn some simple naming for binary covalent compounds.

Sharing is NOT so Equal

Polar vs Non-Polar Bonding

So a covalent bond is based on the sharing of one or more electron pairs between two non-metal atoms. A "perfect" covalent bond means the sharing is perfectly equal - meaning each atom participating in the bond has an equal share of the electrons. This is basically a 50/50 split of electrons. However, this can only happen perfectly when the two atoms have identical electronegativities (EN values). So to be truly perfectly covalent an atom needs to bond with itself, then you are guaranteed that each atom pulls on the electrons the same amount.

So anytime the two atoms are different, there is most likely a difference in EN values and therefore an unequal sharing of the electrons. This is the basis of polar covalent bonds. A polar covalent bond is when one of the atoms gets a bit more of the electrons - technically meaning an unequal sharing of the electron pair. The more electronegative atom will pull the electrons to itself a bit more than the other atom. This leads to a slight partial negative charge (δ–) on the more electronegative atom and a partial positive charge (δ+) on the more electropositive atom. This is a polar bond. Polar bonds can add up on a molecule to give a polar molecule which has a net dipole. Please also note that polar bonds can also just cancel each other out to result in a non-polar molecule as well. You really have to know the three dimensional shape of a molecule plus all the polarities of the bonds to determine if a molecule ends up polar or non-polar. Here is a line/dot structure for HCl (hydrogen chloride) that also shows the partial charges on the H and the Cl.

Chlorine is more electronegative than hydrogen which means it will pull the electrons more toward itself and away from hydrogen. The result is a partial negative charge on the chlorine and a partial positive charge on the hydrogen. Anytime there is a net separation of positive and negative charge for a molecule, the molecule is polar and will have a net dipole moment which is just a measure of the partial charge separation.

Nomenclature for binary (simple) covalent compounds

What I mean by binary covalent compounds is that only two different elements make the compound. It is a bit like ionic compounds except instead of a cation and an anion, you have element1 and element2 and those elements are non-metals.

Who's first? In a similar way to ionic compounds (cations are first), you should always list the more electropositive element first - that is the one with the lower electronegativity value (another reason to learn and memorize the trend).

Prefixes: For covalent compounds we will have to use prefixes to tell others how many of an atom there is in the compound. A very simple example of this is carbon monoxide (CO) and carbon dioxide (CO2). The mono- prefix means one and the di- prefix means two. Can you gues what SF3 is? Sulfur trifluoride. Notice we name the second element (which is always the more electronegative one) as an -ide, just like we did for monatomic anions. Learn your prefixes so you can get the counts right.

mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, and deca- will cover you from 1 to 10 which is plenty.

Avoid Double Vowels (sometimes): When using a prefix ending in o or a for the number (which is all of them except di- and tri-) you might need to drop the o or a when you combine it with oxide. Examples are to write and say carbon monoxide, and do NOT say or write carbon monooxide (note the double vowel). N2O5 is dinitrogen pentoxide (not pentaoxide). This avoid double vowels is mainly for fixing the names of oxides. You DO keep a double vowel for something with iodide. So CI4* is carbon tetraiodide (not tetriodide). So fix it on oxides and avoid the ao and oo double vowels.

* Look out for the san-serif font thing. That is a capital C and a capital i, I. Here's a lowercase ell, l and an uppercase I. Our exams are printed in a serif font so it is obvious: CI4.

A friendly reminder: Do NOT get confused and start using your prefix knowledge with ionic compounds. Remember, ionic compounds just name the ions. Covalent compounds name the elements where the first is the element and the second is the -ide version AND we use prefixes for counts.

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