Explain groups, periods, metals, nonmetals, metalloids, families, and valence patterns. Use table position to predict common ions and reactivity.
A lesson of Chemistry, in Atomic Structure and Periodicity, about 35 minutes.
By the end of this lesson, you will be able to use a group number to predict valence electrons and common ions, name the families and what each one does, tell metals from nonmetals and metalloids, and say why the table's shape is what it is.
Groups Behave Alike Because They Match
A group is a column, and elements in one share the same number of valence
electrons. That is the whole reason the table exists, and it was noticed from
chemical behavior half a century before anyone knew what an electron was.
The period number is the outermost energy level. Period $3$ elements have
their valence electrons in level $3$.
Mendeleev arranged the table by properties in 1869 and left gaps where the
pattern demanded an element nobody had found. He predicted the mass, density and
oxide formula of the missing element below silicon; germanium was discovered
fifteen years later and matched closely. A classification that predicts is doing
more than sorting.
Valence Electrons From the Group Number
For main-group elements, no counting is needed.
Group
Valence electrons
Common ion
$1$
$1$
$+1$
$2$
$2$
$+2$
$13$
$3$
$+3$
$14$
$4$
rarely ionic
$15$
$5$
$-3$
$16$
$6$
$-2$
$17$
$7$
$-1$
$18$
$8$
none
Groups $13$ to $18$: subtract $10$ to get the valence count.
Everything is aiming at eight. Atoms gain, lose or share electrons to reach a
full outer level, which is eight for most elements, and this is the octet
rule. Whether an element loses or gains is decided purely by which is fewer
moves: group $1$ loses one, and group $17$ gains one, and neither would consider
doing it the other way round.
Group $14$ sits exactly halfway, four from either target, which is why carbon
shares rather than transfers and why carbon chemistry is a subject of its own.
Sort It: Which Ion Will It Form
Example 1: Predicting From Position
Here is every step, worked all the way through.
An element in period $3$, group $16$.
How many valence electrons does it have? 6
Which energy level are they in? 3
To reach eight, does it gain or lose electrons? Type one word gain
How many does it gain? 2
What charge does the ion carry, as a number? -2
The Families
Six groups have names, because each behaves distinctively enough to deserve one.
Alkali metals (group 1). One valence electron, given up eagerly. The most
reactive metals there are, and reactivity increases DOWN the group, so cesium is
far more violent with water than lithium. Never found uncombined in nature.
Alkaline earth metals (group 2). Two valence electrons, reactive but less so.
Calcium and magnesium are structural in bone and chlorophyll respectively.
Transition metals (groups 3 to 12). $d$-block. Hard, high-melting, and
distinctive for variable charges: iron forms $\mathrm{Fe^{2+}}$ and
$\mathrm{Fe^{3+}}$, copper $\mathrm{Cu^{+}}$ and $\mathrm{Cu^{2+}}$. Their
compounds are usually colored, and the color comes from $d$ electrons absorbing
visible light.
Halogens (group 17). Seven valence electrons, one short. The most reactive
nonmetals, and reactivity increases UP the group, so fluorine is the most
reactive element in existence. All form $-1$ ions and diatomic molecules.
Noble gases (group 18). A full outer level, so almost no reactivity at all.
Monatomic, unlike every other gas element.
Lanthanides and actinides. The $f$-block, pulled out below to keep the table
a readable width. All actinides are radioactive.
Sort It: Which Family
Metals, Nonmetals and the Staircase
A zigzag line runs down the right side of the table, and it separates two very
different kinds of behavior.
Metals, to the left, are the large majority. Shiny, conductive, malleable,
and they lose electrons to form positive ions. They hold their valence
electrons loosely, which is exactly why they conduct: the electrons move freely
through the whole lattice.
Nonmetals, to the upper right, are dull, brittle when solid, and insulating.
They gain or share electrons to form negative ions.
Metalloids sit on the staircase itself: boron, silicon, germanium, arsenic,
antimony, tellurium. They behave in between, and their most important property is
being semiconductors, conducting under some conditions and not others.
That in-between behavior is not a curiosity. Silicon's ability to be switched
between conducting and insulating is the basis of every transistor, and therefore
of every computer. The most technologically important elements sit exactly on
the boundary, which is a good argument for paying attention to boundary cases.
Worked Example 1: Reading the Position
An element in period $4$, group $17$.
Why the Table Has That Shape
The blocks are the sublevels, and the widths are their capacities.
Period lengths follow from which sublevels are available. Period $1$ has only
$1s$, so it holds two elements. Period $2$ has $2s$ and $2p$, so eight. Period $4$
adds $3d$, so eighteen.
$$2, ; 8, ; 8, ; 18, ; 18, ; 32$$
Those numbers are not arbitrary, and they were a complete mystery until electron
configuration explained them. Mendeleev knew the pattern of repetition and had no
idea why it repeated.
The gap in periods 6 and 7 is where the $f$-block belongs. Inserting fourteen
extra columns would make the table impossibly wide on a page, so those rows are
conventionally printed below. They are part of the table, not an appendix.
Worked Example 2: Period Lengths
Worked Example 3: Predicting a Compound
An element in group $2$ combines with one in group $17$.
Common Mistakes
Mistake 1: Using the group number directly above group 12
Subtract $10$: group $16$ has six valence electrons.
Mistake 2: Expecting the metal to gain electrons
Metals lose; nonmetals gain.
Mistake 3: Assuming reactivity runs the same way in every group
Alkali metals get MORE reactive down; halogens get LESS.
Mistake 4: Giving transition metals one fixed charge
Iron is $\mathrm{Fe^{2+}}$ or $\mathrm{Fe^{3+}}$.
Mistake 5: Treating the $f$-block as separate from the table
It belongs in periods $6$ and $7$; it is printed below for width.
Where This Goes Next
The octet rule is a good approximation with well-known exceptions, and knowing
them is what separates using the rule from believing it.
Hydrogen and helium want two, not eight, because level $1$ only has an $s$
sublevel.
Boron is stable with six in $\mathrm{BF_3}$, which is why it is such an
aggressive electron acceptor.
Elements from period $3$ down can EXCEED eight, because empty $d$ orbitals are
available. $\mathrm{SF_6}$ has twelve electrons around the sulfur, and it is a
perfectly stable industrial gas. That is impossible for period $2$ elements, which
have no $d$ sublevel at all, and it explains why nitrogen cannot form
$\mathrm{NF_5}$ while phosphorus forms $\mathrm{PF_5}$ readily.
The noble gases were themselves an exception to their own reputation. They were
called "inert" until 1962, when $\mathrm{XePtF_6}$ was made. Xenon's outer
electrons are far enough from its nucleus to be pulled away by something
sufficiently aggressive, so several xenon compounds now exist. The lighter noble
gases still resist everything.
The next lesson explains all of this with one idea: effective nuclear charge,
how strongly the nucleus actually pulls on an outer electron once the inner ones
have shielded it. Every trend across and down the table, and every exception here,
comes out of that single quantity.
Check Your Understanding
Five of them, no instructions this time.
Summary
A group is a column with matching valence electrons, so its members behave
alike. A period is a row, changing gradually.
Group number gives valence electrons (subtract $10$ above group $12$) and
therefore the common ion.
Metals lose, nonmetals gain, metalloids sit on the staircase and are
semiconductors.
The block widths are the sublevel capacities, which is why the periods hold
$2, 8, 8, 18, 18, 32$.
Position predicts formula, which is what the table is for.