The unit on one page

Every bond in this unit is one of four kinds, and each kind is defined by what the valence electrons do.

Learning objectives

These are the twelve “I can” targets from the first slide of your Bonding deck. Click one to jump to its section.

The slides also cover two targets that are not on that list: bond energy (breaking is endothermic, forming is exothermic, in section 8) and why ions are more stable in their lattice (in section 5).

TL;DR

• Bonds are about valence electrons. They get transferred (ionic), shared (covalent), or pooled (metallic).
• Location predicts behavior. Metals (left) give electrons away; nonmetals (right) take or share them.
• Electronegativity (EN) is how hard an atom pulls on shared electrons. Subtract: ΔEN ≤ 0.4 nonpolar covalent, 0.4 to 1.8 polar covalent, above 1.8 ionic.
• Lewis structures show bonding pairs and lone pairs so every atom gets 8 (octet), or 2 for hydrogen (duet). More shared pairs = shorter, stronger bonds.
• Breaking a bond absorbs energy (endothermic). Forming a bond releases energy (exothermic).
• Shape (VSEPR) decides polarity. A molecule is polar only if its bond dipoles do not cancel.
• The electron sea explains why metals conduct, bend, and shine.

Glossary

Valence electrons

The electrons in an atom’s outermost shell. They are the only electrons involved in bonding. For main-group elements, count them from the group number (Group 1 has 1, Group 17 has 7).

Chemical bond

An attraction that holds atoms or ions together. Atoms bond because the bonded arrangement has lower energy, so it is more stable.

Octet rule

Atoms tend to gain, lose, or share electrons until they have 8 valence electrons, like a noble gas.

Duet rule

Hydrogen (and helium) is stable with only 2 electrons. Hydrogen always forms exactly one bond.

Ionic bond

The attraction between oppositely charged ions after electrons are transferred from a metal to a nonmetal. Example: NaCl.

Cation and anion

A cation is a positive ion (it lost electrons; metals form cations). An anion is a negative ion (it gained electrons; nonmetals form anions).

Crystal lattice

The repeating 3D pattern of ions in an ionic solid. Each ion is surrounded by ions of the opposite charge.

Covalent bond

A bond in which two nonmetal atoms share a pair of electrons. Example: the H–H bond in H₂.

Polar covalent bond

A covalent bond with unequal sharing. The more electronegative atom gets a partial negative charge (δ−), the other a partial positive charge (δ+). ΔEN 0.4 to 1.8.

Nonpolar covalent bond

A covalent bond with equal (or nearly equal) sharing, so there are no partial charges. ΔEN 0.4 or less.

Metallic bond

The attraction between metal cations and a sea of delocalized valence electrons that flows around them.

Delocalized electrons

Electrons that are not tied to one atom or one bond. In a metal they belong to the whole piece of metal.

Electronegativity (EN)

The ability of an atom to attract the shared electrons in a bond. It rises across a period and up a group; fluorine is highest (4.0).

ΔEN (electronegativity difference)

The larger EN minus the smaller EN for the two bonded atoms. It decides the bond type.

Lewis structure

A drawing of a molecule that shows every valence electron: lines for bonding pairs, dots for lone pairs.

Bonding pair and lone pair

A bonding pair is shared between two atoms (one line). A lone pair belongs to one atom and is not shared (two dots).

Single, double, triple bond

One, two, or three shared pairs (2, 4, or 6 shared electrons). More pairs make the bond shorter and stronger.

Bond length

The distance between the nuclei of two bonded atoms, at the bottom of the energy well.

Bond energy

The energy needed to break one mole of a bond (kJ/mol). The same amount is released when that bond forms.

Endothermic and exothermic

Endothermic processes absorb energy (bond breaking). Exothermic processes release energy (bond forming).

Dipole

A separation of charge: a positive end and a negative end. A bond dipole arrow points toward the δ− atom.

VSEPR theory

Valence Shell Electron Pair Repulsion: electron domains around a central atom spread as far apart as possible, which sets the molecule’s shape.

Electron domain

Any region of electrons around the central atom: a single, double, or triple bond, or a lone pair, each counts as one.

Malleable, ductile, lustrous

Malleable: can be hammered into sheets. Ductile: can be drawn into wire. Lustrous: shiny. All three come from the electron sea.

Alloy

A mixture of a metal with other elements (brass is copper + zinc; steel is iron + a little carbon).

Chromatography

A way to separate a mixture: a solvent (mobile phase) carries the parts across paper (stationary phase) at different speeds.

Rf value

Retention factor: distance a substance moved ÷ distance the solvent front moved. Always between 0 and 1.

1. What electrons can do

Target #1. Before anything else, a prediction from your first slide.

PREDICT: describe the behavior of electrons when atoms bond. Which is closest to your idea?

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The answer is all three. Only the valence electrons (the outermost ones) take part, and what they do depends on which atoms are bonding. Watch each version happen.

Transfer, share, or pool

⏸ Pause ↺ Restart

A metal meets a nonmetal. Sodium has 1 valence electron; chlorine has 7 and wants 8.

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Here is the class comparison from the slides, extended with metals.

Your slides ask three PREDICT questions. Answer each in your head, then open it.

Q PREDICT 1: Does a higher boiling point or a lower boiling point take more energy?

A higher boiling point takes more energy. Boiling point is the temperature where a liquid turns into a gas, so the particles must be pulled apart. The stronger the attractions holding them together, the more heat (energy) you need, and the higher the boiling point.

Q PREDICT 2: Do ionic compounds or covalent molecules have higher boiling points?

Ionic compounds. To boil salt you must pull apart ions held by strong + to − attractions in every direction. To boil a covalent substance you only separate whole molecules from each other, and those attractions between molecules are much weaker. (The covalent bonds inside each molecule do not break when it boils.)

Q PREDICT 3: Which is more likely to be a solid at room temperature, ionic or covalent?

Ionic. Room temperature (about 25 °C) is far below the melting point of nearly every ionic compound, so they are solid crystals. Many covalent substances melt or boil below 25 °C, so they are liquids or gases. Your slide’s follow-up, “can we predict which polar covalent substances are liquids and which are gases?”, is a preview of intermolecular forces: more polar and heavier molecules attract each other more, so they tend to be liquids instead of gases.

Now check the predictions against real data. Each dot is one substance; the dashed line is room temperature, so anything to its left is already melted at 25 °C.

Hover a dot for its name and melting point. Ionic and metallic substances cluster far right; most covalent substances sit left of room temperature. The exceptions (iodine, sugar, mercury, sodium) are why we say 'usually'. CO₂ turns straight from solid to gas at −78 °C. Table view: the same data is listed in the Cheat sheet.

Want the story in six minutes? This TED-Ed video is the “extra resource” from your slides.

2. Who gives and who takes: the periodic table

Target #3. You can predict what an atom will do with its electrons just from where it sits.

• Metals (left of the staircase) have only 1 to 3 valence electrons. Losing them is easier than gaining five or more, so metals give electrons away and become cations (+).
• Nonmetals (right of the staircase, plus hydrogen) are 1 to 3 electrons short of an octet. They take electrons from metals (becoming anions, −) or share with other nonmetals.
• Noble gases (Group 18) already have a full shell, so they almost never bond.
• Metalloids (the staircase: B, Si, Ge, As, Sb, Te, and Po on your slide) are in between.

That gives three pairing rules: metal + nonmetal → transfer (ionic), nonmetal + nonmetal → share (covalent), metal + metal → pool (metallic). Try it. Click two elements, or use a quick pick.

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Who gives, who takes?

Click any two elements. Click the same one twice for a bond between identical atoms.

Metal or nonmetalElectronegativityValence electrons

H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og La–Lu Ac–No

Bond Atom 1 + Atom 2 Clear Try: Na–ClMg–OC–OH–HCu–ZnAl–Cl

Use the shortcut as a key: answer one question at a time.

K

Classify a bond from its formula

Click to answer. Each answer narrows the identification.

Step 1 Start over

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Let's practice: from the slides

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1. 01

   CaF₂ (calcium fluoride)

   A Ionic: a metal plus a nonmetal B Covalent: two nonmetals sharing C Metallic: two metals pooling D No bond: a noble gas is present
   Why: Ca is a metal (Group 2) and F is a nonmetal, so electrons transfer: Ca²⁺ and two F⁻.
2. 02

   CO₂ (carbon dioxide)

   A Ionic: a metal plus a nonmetal B Covalent: two nonmetals sharing C Metallic: two metals pooling D No bond: a noble gas is present
   Why: C and O are both nonmetals, so they share electrons. ΔEN = 3.5 − 2.5 = 1.0, so each C=O bond is polar covalent.
3. 03

   NaCl (sodium chloride)

   A Covalent: two nonmetals sharing B Metallic: two metals pooling C Ionic: a metal plus a nonmetal D No bond: a noble gas is present
   Why: Na is a metal, Cl a nonmetal. ΔEN = 3.0 − 0.9 = 2.1, above 1.8: ionic.
4. 04

   CH₄ (methane)

   A Ionic: a metal plus a nonmetal B Metallic: two metals pooling C No bond: a noble gas is present D Covalent: two nonmetals sharing
   Why: C and H are both nonmetals. ΔEN = 2.5 − 2.1 = 0.4, which your class calls nonpolar covalent.
5. 05

   Br₂ (bromine)

   A Covalent: two nonmetals sharing B Ionic: a metal plus a nonmetal C Metallic: two metals pooling D No bond: a noble gas is present
   Why: Two identical nonmetal atoms share a pair exactly equally: nonpolar covalent, ΔEN = 0.

3. Electronegativity: the tug-of-war number

Target #2. Electronegativity is the ability of an atom to attract the shared electrons in a bond. Think of a bond as a tug-of-war over the electron pair; EN is each atom’s strength.

• It increases across a period (left to right): more protons pull on the same shell.
• It increases up a group (bottom to top): smaller atoms hold electrons closer to the nucleus.
• Fluorine (4.0) is the strongest; cesium and francium (0.7) are the weakest. Most noble gases have no value.

Your slides show the main-group values as a 3D skyline. Here it is, spinnable: drag to rotate.

Bar height = electronegativity; hover a bar for its value. The skyline climbs toward fluorine in the back right corner. Drag to rotate, scroll to zoom. Values match the electronegativity table on your slides.

The ΔEN rule

To classify a bond, subtract the smaller EN from the larger one. That difference, ΔEN, tells you how lopsided the tug-of-war is.

Now use the real numbers. The calculator below does the subtraction, places the bond on the scale, and draws what the electrons do.

⚛

ΔEN calculator

Click any two elements. Click the same one twice for a bond between identical atoms.

Metal or nonmetalElectronegativityValence electrons

H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og La–Lu Ac–No

Bond Atom 1 + Atom 2 Clear Try: H–HB–HC–HH–ClO–HNa–ClH–FMg–I

The four worked examples from your slides:

Q

ΔEN check

0 of 4 answered

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1. 01

   What is ΔEN for the N–H bond (N = 3.0, H = 2.1), and what bond type is it?

   A 0.9, polar covalent B 0.9, nonpolar covalent C 5.1, ionic D 2.1, ionic
   Why: 3.0 − 2.1 = 0.9, which falls between 0.4 and 1.8: polar covalent. (Adding the values is a common slip.)
2. 02

   Which bond is the MOST polar?

   A C–H (2.5 and 2.1) B C–Cl (2.5 and 3.0) C O–H (3.5 and 2.1) D S–H (2.5 and 2.1)
   Why: O–H has the largest difference, 1.4. C–Cl is 0.5, C–H and S–H are 0.4.
3. 03

   KF has ΔEN = 4.0 − 0.8 = 3.2. Which picture matches its electrons?

   A A pair shared equally in the middle B A pair shared but pulled toward F C Electrons spread in a sea around ions D An electron moved fully from K to F
   Why: ΔEN is far above 1.8, so the electron transfers: K⁺ and F⁻.
4. 04

   Electronegativity generally increases as you move

   A down a group and to the left B up a group and to the right C down a group and to the right D up a group and to the left
   Why: EN rises toward the top right corner of the table, peaking at fluorine (noble gases aside).

4. Polar vs nonpolar bonds up close

Targets #1 and #2. Your slides define these from pictures. Here is the same idea as electron clouds: where the shared electrons spend their time.

The δ symbol (Greek “delta”) means partial. A polar bond has partial charges that are less than a full +1 or −1, because the electron is still shared, just unevenly. An ionic bond has full charges because the electron moved all the way over.

Match the bond to what its electrons do 0 / 5 matched

Terms

F–FC–HO–HLiFCu–Cu

Definitions

Electron transferred; Li⁺ and F⁻ formOne pair shared perfectly equallyElectrons pooled in a sea around cationsPair shared, but pulled toward one atomShared almost equally (ΔEN 0.4)

5. Ionic compounds and the crystal lattice

Slide target: compare the stability of ions when they are separated vs. in their lattice. An ionic compound is not a molecule. It is a giant crystal lattice: a 3D checkerboard where every cation is surrounded by anions and every anion by cations. Spin the real arrangement of table salt.

3D

Crystal lattices in 3D

Drag to rotate, scroll or pinch to zoom.

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Your slides list four properties of ionic compounds. Each one comes straight from the lattice.

Why ionic compounds dissolve in water

Water is polar: the oxygen end is slightly negative (δ−) and the hydrogen ends are slightly positive (δ+). Those ends grab the ions on the surface of the crystal.

Dissolving salt in water

⏸ Pause ↺ Restart

A salt crystal drops into water. Na⁺ and Cl⁻ are locked in the lattice.

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When do ions conduct?

Electricity is moving charge. In an ionic compound the charges are the ions, so the question is always: can the ions move? Click through the four tests from your slide.

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Your slide’s Turn and Talk questions:

Q Why do ionic bonds have high melting points? Have you ever been able to melt salt?

Every ion is attracted to all the opposite charges around it, so melting means overcoming a huge number of strong attractions. A kitchen stove tops out around 250 to 300 °C; salt needs 801 °C, which is why you have never melted it. (Sugar, which is covalent, caramelizes on a stove at about 186 °C.)

Q How can ionic compounds conduct electricity easily?

Only when the ions are free to move: dissolved in water or melted. A solid crystal does not conduct because its ions are fixed in place.

Q PREDICT: do covalent compounds dissolve and conduct electricity too?

Some dissolve, but they do not conduct. Polar covalent substances like sugar dissolve in water, but the molecules stay whole, so no ions form and the bulb stays dark. Nonpolar covalent substances like wax do not even dissolve in water (“like dissolves like”). You will test exactly this in the bonding lab.

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Ionic compounds check

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1. 01

   Why does solid NaCl NOT conduct electricity?

   A It has no charged particles in it B Its ions are fixed in the lattice C Its electrons are shared equally D It is made of separate molecules
   Why: The ions are there, but locked in place. Conduction needs moving charges, which happens once the lattice is dissolved or melted.
2. 02

   When NaCl dissolves, which end of a water molecule faces a Na⁺ ion?

   A The oxygen end, which is δ− B The hydrogen end, which is δ+ C Either end, it does not matter D Neither, water ignores ions
   Why: Opposites attract: the partially negative oxygen points at the positive sodium ion.
3. 03

   The best description of solid sodium chloride is

   A a single NaCl molecule B pairs of NaCl molecules C a lattice of Na⁺ and Cl⁻ ions D Na atoms sharing with Cl atoms
   Why: NaCl is a giant lattice. The formula only gives the 1:1 ratio of ions.
4. 04

   Why are ions more stable in a lattice than when separated?

   A Ions in a lattice have lost their charge B The lattice shares electrons evenly C Lattice ions turn back into neutral atoms D Each ion is attracted to many opposite charges
   Why: Many simultaneous + to − attractions lower the energy. Lower energy = more stable.

6. Lewis structures: the duet and octet rules

Target #4. A Lewis structure shows every valence electron in a molecule. Lines are bonding pairs (shared), and dots are lone pairs (not shared). The goal: every atom ends up with 8 electrons around it (the octet rule), except hydrogen, which needs only 2 (the duet rule).

First you need to count valence electrons. For main-group elements it is the last digit of the group number. Flip the table to “Valence electrons”.

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Valence electrons at a glance

Hover or tap an element for details.

Valence electronsMetal or nonmetal

H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og La–Lu Ac–No

Each atom’s valence electrons can be drawn as a Lewis dot symbol: one dot per electron, up to two per side.

Unpaired dots are the electrons available for bonding. Carbon has 4 unpaired, so it makes 4 bonds; nitrogen 3; oxygen 2; fluorine 1; neon 0. Hydrogen, with 1, always makes exactly 1 bond.

The six-step method

Work through ammonia, NH₃.

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When there are not enough electrons to give everyone an octet with single bonds, atoms share more pairs. Carbon dioxide shows how.

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Ionic compounds get Lewis structures too, but nothing is shared: show the transfer with charges and brackets.

The classic covalent structures you should be able to draw cold:

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Count the valence electrons

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Check all Reset

7. Single, double, and triple bonds

Target #6. A single bond shares 1 pair (2 electrons), a double bond shares 2 pairs (4 electrons), and a triple bond shares 3 pairs (6 electrons). More shared electrons pull the two nuclei closer together and hold them tighter.

More shared pairs → shorter bond → stronger bond. Feel it: drag the atoms apart and watch the energy. The bottom of the well is the bond length, and its depth is the bond energy. Then tick “Compare carbon bonds”.

E

Pull a bond apart

Drag the slider to change the distance between the two nuclei.

Distance

Break the bond Form the bond Compare carbon bonds

Rank these bonds from STRONGEST to weakest (use the chart)

Shuffle Check

• N–N in hydrazine
• C≡C in ethyne
• O=O in oxygen gas
• C=C in ethene
• N≡N in nitrogen gas
• C–C in ethane

Your slide’s self-check, worked:

Q Slide self-check: rank CH₄, C₂H₄, and N₂ by bond length and by bond energy

Longest to shortest bond: CH₄ (single) → C₂H₄ (double) → N₂ (triple). Strongest to weakest: N₂ → C₂H₄ → CH₄. The reasoning your class uses is bond order: single bonds are the longest and weakest, triples the shortest and strongest. (Fine print in Common pitfalls: this comparison is cleanest when the same two elements are bonded.)

8. Bond energy: breaking costs, forming pays

Slide target: define bond energy and recognize that bond breaking is endothermic and bond forming is exothermic.

Bond energy is the energy needed to break one mole of a particular bond, in kJ/mol. Bonded atoms sit at the bottom of an energy well (you just played with it), so:

A chemical reaction does both: it breaks the bonds in the reactants and forms the bonds in the products. Whichever is bigger wins. Watch the energy ledger for hydrogen gas burning in chlorine.

▶ Replay

Your slide’s self-check uses a reaction energy diagram. Find each part. (The image is in quiz mode: it names a part and you click the dot.)

Q Slide self-check: is the graph showing the breaking of bonds or the formation of bonds?

Overall, formation. The products end up lower than the reactants, so energy was released: the reaction is exothermic, which means more energy came out of forming new bonds than went into breaking the old ones. Inside the curve you can see both steps: the climb up to the transition state (Eₐ) is energy going in to start breaking bonds, and the long drop down to the products is energy coming out as new bonds form.

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Bond energy check

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1. 01

   Breaking the H–H bond in H₂ requires 436 kJ/mol. What happens when 1 mol of H–H bonds forms?

   A 436 kJ of energy is absorbed B No energy change at all C 436 kJ of energy is released D 872 kJ of energy is released
   Why: Forming a bond releases exactly the energy it takes to break it: 436 kJ out.
2. 02

   A reaction absorbs 900 kJ breaking bonds and releases 1,100 kJ forming bonds. The reaction is

   A exothermic, ΔH = −200 kJ B endothermic, ΔH = +200 kJ C exothermic, ΔH = +2,000 kJ D endothermic, ΔH = −200 kJ
   Why: ΔH = 900 − 1,100 = −200 kJ. More energy out than in: exothermic.
3. 03

   Which bond needs the most energy to break?

   A C–C, a single bond B C=C, a double bond C N–N, a single bond D N≡N, a triple bond
   Why: N≡N (945 kJ/mol) is the strongest on the chart: three shared pairs.
4. 04

   Bond breaking is always

   A exothermic, because energy comes out B endothermic, because energy goes in C exothermic, because the atoms get hot D neither, energy is not involved
   Why: You always have to put energy in to separate bonded atoms. Remember BARF: Breaking Absorbs, Releasing Forms.

9. VSEPR: predicting the shape

Target #12. Lewis structures are flat, but molecules are 3D. VSEPR (Valence Shell Electron Pair Repulsion) says: the electron groups around a central atom repel each other and spread as far apart as possible. Count them, and the shape follows.

How to count electron domains on the central atom: every bonded atom counts as 1 (a single, double, or triple bond is still one domain), and every lone pair counts as 1.

Lone pairs take up more room than bonding pairs, so they squeeze the bonds closer together. That is why the angle shrinks from 109.5° (CH₄) to 107° (NH₃, one lone pair) to 104.5° (H₂O, two lone pairs). Spin each one.

3D

VSEPR shapes in 3D

Drag to rotate, scroll or pinch to zoom.

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M

Match each molecule to its shape

Build your own: add bonds and lone pairs and watch them push apart. Switch to “Real Molecules” to compare.

Sim by PhET, University of Colorado Boulder

10. Polar molecules: add up the dipoles

Target #5. A bond dipole is drawn as an arrow with a plus-sign tail (the δ+ end) pointing toward the more electronegative atom (the δ− end). A molecule is polar only if those arrows do not cancel when you add them up in 3D.

Use the two-question test on any molecule.

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Is the molecule polar?

Click to answer. Each answer narrows the identification.

Step 1 Start over

The polarity slide again, now with the real shapes:

Test any molecule with an electric field: polar molecules turn to line up with it, nonpolar ones do not.

Sim by PhET, University of Colorado Boulder

Q

Molecular polarity check

0 of 4 answered

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1. 01

   CO₂ has two polar C=O bonds. Why is the molecule nonpolar?

   A Its C=O bonds are not really polar B Its linear shape makes the pulls cancel C Carbon has two lone pairs on it D Oxygen is less electronegative than C
   Why: The two dipoles are equal and point in exactly opposite directions (180°), so the net dipole is zero.
2. 02

   Which molecule is POLAR?

   A CCl₄, tetrahedral B BCl₃, trigonal planar C N₂, linear D PCl₃, trigonal pyramidal
   Why: PCl₃ has a lone pair on P, which makes it pyramidal and lopsided. The other three are symmetric or have no polar bonds.
3. 03

   In a bond dipole arrow, the arrowhead points toward

   A the more electronegative atom B the less electronegative atom C the larger atom, whatever its EN D the central atom, always
   Why: The arrow points to the δ− end, the atom that pulls the shared electrons harder.
4. 04

   Water is polar mainly because

   A it is made of hydrogen and oxygen B it is a liquid at room temperature C it is bent, so its O–H dipoles add D its bonds are ionic, not covalent
   Why: Two polar O–H bonds plus a bent shape (from oxygen's two lone pairs) means the dipoles add to a net dipole toward O.

11. Metallic bonding: the electron sea

Targets #1 and #10. In a metal, each atom releases its valence electrons. The electrons become delocalized: they belong to the whole piece of metal, not to any one atom. What is left is a rigid framework of cations immersed in a sea of electrons that acts as the cement holding the 3D network together. That electrostatic attraction between the cations and the sea is the metallic bond.

Every metal property in your slides falls out of that one picture. Try all four tabs.

e⁻

Inside a metal

Electron seaAdd a voltageHeat one endHit it with a hammer

High melting points. Generally high melting points show that metallic bonding is strong. Your slide adds: melting points increase with the number of valence electrons each atom adds to the sea, because more electrons (and higher-charged cations) means a stronger attraction.

Na gives 1 electron per atom to the sea, Mg gives 2, Al gives 3. (The Mg to Al step is small; real melting points also depend on how the atoms pack.)

Alloys

An alloy mixes a metal with other elements while keeping the electron sea. There are two ways the extra atoms fit in, as your last slide shows.

The 3D lattice viewer in section 5 has a Copper metal button: compare its tight packing and electron sea with salt’s checkerboard.

Q

Metallic bonding check

0 of 4 answered

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1. 01

   What holds a metal together?

   A Shared pairs between two atoms B Cations attracted to an electron sea C Anions attracted to free protons D Transferred electrons between ions
   Why: The positive metal ions are held in place by their attraction to the delocalized electrons around them.
2. 02

   Why can a copper wire be bent without breaking?

   A The sea holds the layers as they slide B Copper ions repel each other strongly C Copper atoms share fixed electron pairs D Bending melts the copper just a little
   Why: Metallic bonds are non-directional, so a sliding layer stays bonded to the sea around it.
3. 03

   Which metal should have the highest melting point?

   A Sodium, 1 valence electron B Potassium, 1 valence electron C Cesium, 1 valence electron D Aluminum, 3 valence electrons
   Why: More valence electrons in the sea means a stronger metallic bond. Aluminum melts at 660 °C; sodium at 98 °C.
4. 04

   A metal conducts electricity because

   A its ions move toward the electrodes B it dissolves in water to form ions C its electrons are free to move D its atoms share pairs equally
   Why: Mobile delocalized electrons flow toward the + end when a voltage is applied.

12. Conductors and insulators

Target #11. In class you test which solids conduct and rank them. Here is how real materials compare. The scale is logarithmic: each tick is 10 times the one before, and the bars span 23 powers of ten.

Each dot is one material: typical room-temperature values in siemens per meter (S/m). Graphite is the famous nonmetal exception: its carbon layers have delocalized electrons. Glass and rubber vary by type; typical values shown.

Rank from BEST conductor to worst

Shuffle Check

• Iron
• Silver
• Graphite (pencil core)
• Aluminum
• Rubber
• Copper
• Silicon

13. Chromatography: separating a mixture

Target #7. Paper chromatography separates the substances in a mixture (like the dyes in a marker) using what you just learned about polarity: like dissolves like.

• The stationary phase is the paper, which stays put.
• The mobile phase is the solvent, which climbs the paper by capillary action.
• A substance that dissolves well in the solvent rides along and travels far. A substance more attracted to the paper lags behind. Different attractions mean the mixture spreads into separate spots.

Run the virtual strip. Try the permanent marker in both solvents.

Rf

Separate the dyes

Choose an ink and a solvent, then run the strip.

Ink sample

Solvent

Run the strip New strip

Each spot is described by its Rf value (retention factor):

Q Worked example: a blue dye moves 3.2 cm while the solvent moves 8.0 cm. What is its Rf?

Rf = 3.2 cm ÷ 8.0 cm = 0.40. A yellow dye on the same strip at 6.8 cm has Rf = 6.8 ÷ 8.0 = 0.85, so the yellow dye dissolves better in this solvent than the blue one does. The same dye in the same solvent always gives about the same Rf, which is how chromatography identifies substances (forensics uses this to match inks).

Q

Chromatography check

0 of 4 answered

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1. 01

   Why is the start line drawn in pencil, not pen?

   A Pencil is easier to see on wet paper B Pen ink would dissolve and run too C Pencil speeds up the solvent D Pen would block the capillary action
   Why: Graphite does not dissolve in the solvent, so the line stays put. Pen ink is itself a mixture that would separate and confuse your results.
2. 02

   A dye has Rf = 0.9 in water. What does that tell you?

   A It is very soluble in water B It sticks tightly to the paper C It is nonpolar like oil D It did not move from the line
   Why: It traveled almost as far as the water itself, so it dissolves well in water (it is polar).
3. 03

   Permanent marker ink stays on the start line in water but separates in alcohol. Its dyes are probably

   A ionic, like table salt B very polar, like sugar C metallic, like copper wire D mostly nonpolar, like oil
   Why: Like dissolves like: nonpolar dyes will not dissolve in very polar water, but they do dissolve in less polar alcohol.
4. 04

   The solvent level in the beaker must start

   A above the start line B level with the start line C below the start line D at the top of the paper
   Why: If the solvent covers the ink spot, the ink dissolves into the beaker instead of climbing the paper.

14. The bonding lab: identify the unknowns

Targets #8 and #9. From your lab slide: you will be given two or three unknown substances; use what you have learned about the properties of ionic and covalent bonds to determine what the substances are, by designing an investigation based on melting point, solubility, and conductivity.

First, know the fingerprint each bond type leaves in the data.

The two tests that do the most work, as a decision tree:

Now run the lab. Every time you click “New unknowns”, the samples are reshuffled behind new letters.

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Virtual lab: six unknown solids

Pick a sample, run tests, then classify every sample from your data.

New unknowns

1 · Sample shelf

2 · Run a test on sample A

3 · Data table (fills in as you test)

Sample	Hot plate	In water	Solid conducts?	Solution conducts?	Hammer	4 · Your call

Check my answers

Writing the lab report

Your group report must include every part on this checklist (from the lab slide).

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Practice the conclusion (claim, evidence, reasoning)

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Worked example

Put every target together on one molecule you have not seen yet: phosphorus trichloride, PCl₃.

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The same workflow for an ionic compound is shorter, because there is no shape or polarity to find.

Practice

Try each one on paper first, then open it.

Q 1. Classify each bond using ΔEN: (a) C–S (b) Li–Cl (c) P–H (d) O–F (e) Ca–O

(a) C–S: 2.5 − 2.5 = 0.0, nonpolar covalent. (b) Li–Cl: 3.0 − 1.0 = 2.0, ionic. (c) P–H: 2.2 − 2.1 = 0.1, nonpolar covalent. (d) O–F: 4.0 − 3.5 = 0.5, polar covalent (F is δ−). (e) Ca–O: 3.5 − 1.0 = 2.5, ionic.

Q 2. Which element in each pair is more likely to LOSE electrons: (a) Na or Cl (b) Mg or S (c) K or Br?

The metal every time: Na, Mg, K. They sit on the left of the table, have 1 or 2 valence electrons, and low electronegativity, so they form cations. The nonmetals (Cl, S, Br) gain electrons to become anions.

Q 3. Draw the Lewis structure of HCN.

10 valence electrons (1 + 4 + 5). C goes in the middle (H is never central). H–C uses 2 and a C≡N triple bond uses 6; the last 2 are a lone pair on N. H has 2, C has 8, N has 8.

1 single bond, 1 triple bond, 1 lone pair

Q 4. Draw the Lewis structure of formaldehyde, CH₂O.

12 valence electrons (4 + 2 + 6). C is central. With only single bonds C would have 6, so one O lone pair becomes a second bond: C=O, plus two C–H bonds and two lone pairs on O.

Trigonal planar around C; polar (the C=O dipole is not canceled)

Q 5. Draw OF₂ and predict its shape and polarity.

20 valence electrons (6 + 2 × 7). O is central with 2 bonds and 2 lone pairs: 4 domains, so bent (like water). O–F is polar (ΔEN 0.5) and the bent shape keeps the dipoles from canceling, so OF₂ is polar.

Bent, polar

Q 6. CCl₄ has four polar C–Cl bonds. Is the molecule polar?

No. Carbon has 4 domains and no lone pairs, so CCl₄ is a perfect tetrahedron with four identical outer atoms. The four equal dipoles cancel in 3D, so the molecule is nonpolar, even though every bond is polar.

Q 7. Use bond energies to find ΔH for N₂ + 3 H₂ → 2 NH₃ (N≡N 945, H–H 436, N–H 391 kJ/mol).

Break: 1 N≡N + 3 H–H = 945 + 3(436) = 2,253 kJ in. Form: 2 NH₃ have 6 N–H bonds = 6(391) = 2,346 kJ out. ΔH = 2,253 − 2,346 = −93 kJ: exothermic.

Q 8. A red dye travels 4.5 cm while the solvent front travels 7.5 cm. What is the Rf? Is the dye more attracted to the solvent or the paper?

Rf = 4.5 ÷ 7.5 = 0.60. It moved more than halfway with the solvent, so it is somewhat more attracted to the solvent than to the paper. A dye with Rf 0.2 on the same strip would cling to the paper more.

Q 9. Lab data: Sample Q melts on the hot plate within seconds, does not dissolve in water, and does not conduct. Sample R does not melt, does not dissolve, conducts as a solid, and flattens under a hammer. Identify both.

Q: nonpolar covalent (very low melting point and insoluble in polar water: wax-like). R: metallic (conducts as a solid and is malleable: electron sea).

Q 10. Explain, using the electron sea, why aluminum foil can be folded but a salt crystal shatters.

In aluminum the cations sit in a sea of delocalized electrons that attracts them in every direction, so layers slide over each other and stay bonded. In a salt crystal, sliding a layer lines up + next to + and − next to −; those like charges repel and the crystal splits apart.

Quiz

Q

Bonding unit test prep

0 of 12 answered

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1. 01

   Which pair of elements forms an ionic bond?

   A K and Br B C and O C Cu and Zn D N and H
   Why: K is a metal and Br a nonmetal (ΔEN 2.0). C–O and N–H are covalent; Cu–Zn is metallic (brass).
2. 02

   In a nonpolar covalent bond, the electrons are

   A transferred from one atom to the other B shared equally between the two atoms C pooled in a sea around many cations D shared but pulled toward one atom
   Why: Nonpolar = equal sharing, because the two atoms pull equally (ΔEN 0.4 or less).
3. 03

   Mg has EN 1.2 and O has EN 3.5. The Mg–O bond has

   A ΔEN 1.3, so it is polar covalent B ΔEN 2.3, so it is polar covalent C ΔEN 2.3, so it is ionic D ΔEN 4.7, so it is ionic
   Why: 3.5 − 1.2 = 2.3, above 1.8: ionic. Also a metal with a nonmetal.
4. 04

   How many valence electrons are in one molecule of HCN?

   A 8 B 12 C 14 D 10
   Why: H 1 + C 4 + N 5 = 10.
5. 05

   Compared with a C–C single bond, a C≡C triple bond is

   A shorter and stronger B longer and stronger C shorter and weaker D longer and weaker
   Why: Three shared pairs pull the nuclei closer (120 vs 154 pm) and hold them tighter (839 vs 347 kJ/mol).
6. 06

   Which statement about breaking a chemical bond is correct?

   A It releases energy, so it is exothermic B It absorbs energy, so it is endothermic C It releases energy, so it is endothermic D It absorbs energy, so it is exothermic
   Why: Breaking always needs energy input: endothermic. (BARF: Breaking Absorbs, Releasing Forms.)
7. 07

   What is the shape of NH₃?

   A trigonal planar B tetrahedral C bent D trigonal pyramidal
   Why: 4 domains with 1 lone pair: trigonal pyramidal, about 107°.
8. 08

   Which molecule is nonpolar?

   A H₂O B NH₃ C CO₂ D HCl
   Why: CO₂ is linear, so its two C=O dipoles cancel. The others have uncanceled dipoles.
9. 09

   Metals conduct heat well because

   A their ions are free to move around B their mobile electrons carry kinetic energy C their atoms share fixed electron pairs D their lattices dissolve when heated
   Why: Delocalized electrons pick up energy in the hot region and carry it quickly through the lattice.
10. 10

    An unknown solid conducts electricity as a solid and bends instead of breaking. It is most likely

    A metallic B ionic C polar covalent D nonpolar covalent
    Why: Only metals conduct as solids (electron sea), and only metals are malleable.
11. 11

    A dye moves 2.0 cm and the solvent front moves 8.0 cm. The dye's Rf is

    A 4.0 B 0.40 C 2.5 D 0.25
    Why: Rf = 2.0 ÷ 8.0 = 0.25. Rf is always between 0 and 1.
12. 12

    Salt water conducts electricity but sugar water does not, because

    A salt is more soluble than sugar B sugar molecules are too big to move C salt releases ions; sugar stays as molecules D sugar water is colder than salt water
    Why: Conduction needs moving charges. Dissolved salt is Na⁺ and Cl⁻ ions; dissolved sugar is whole, neutral molecules.

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Valence electrons

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Outermost electrons, the only ones used in bonding. Main group: last digit of the group number.

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Bonding vocabulary Check

Mnemonics

Common pitfalls

The covalent bond itself is strong (C–C needs 347 kJ/mol, N≡N 945). What is weak are the attractions between separate molecules. When water boils, no O–H bonds break; the molecules just pull away from each other. That is what the slide’s “weak bonds (comparatively)” and “low boiling points” are really about.

Ionic compounds form a lattice, not molecules. “NaCl” only says there is one Na⁺ for every Cl⁻. (Some diagrams label the product “ionic molecule”; on a test, say ionic compound or formula unit.)

Only if the dipoles do not cancel. CO₂, BCl₃, and CCl₄ all have polar bonds and are nonpolar because their shapes are symmetric. Always check the shape.

The slide’s IF₃ drawing leaves out iodine’s two lone pairs. Iodine has 5 electron domains (3 bonds + 2 lone pairs), so the molecule is T-shaped and lopsided: by the VSEPR rules it is polar. If it comes up on a test, ask your teacher which answer they want; the chemistry says polar. (This is also an octet exception: iodine holds 10 electrons.)

Bond order only predicts length when the same two elements are bonded: C–C (154 pm) is longer than C=C (134 pm). Across different elements, atom size matters too: the C–H bonds in CH₄ are actually only 109 pm, shorter than the C=C in C₂H₄, because hydrogen is tiny. Your slide’s self-check ranks by bond order (CH₄ > C₂H₄ > N₂), which is the reasoning your class is testing.

Most do, but not all. Mercury melts at −39 °C (it is a liquid), gallium at 30 °C (it melts in your hand), and sodium at 98 °C. Say “metals usually have high melting points.”

“Occurs between metals” means between metal atoms. A pure copper wire, with only one element, is held together by metallic bonding. Alloys (brass, steel) are just one kind of metallic solid.

Solid ionic compounds do not conduct: their ions are locked in the lattice. They conduct only when dissolved or molten. And dissolving alone is not proof of ionic: sugar dissolves but never conducts.

Breaking a bond always absorbs energy. When a reaction gives off energy, it is because the new bonds that form release more than the old ones cost to break.

A double or triple bond is one domain (it points in one direction). CO₂ has 2 domains on carbon, not 4, which is why it is linear and not tetrahedral.

Cheat sheet

Q Table view: the melting-point chart from section 1

Ionic: MgO 2852 °C · CaF₂ 1418 · NaF 993 · LiF 845 · NaCl 801 · CaCl₂ 772 · KCl 770. Metallic: W 3422 · Fe 1538 · Cu 1085 · Au 1064 · Al 660 · Na 98 · Hg −39. Polar covalent: sucrose 186 (breaks down) · glucose 146 · water 0 · ammonia −78 · HCl −114 · ethanol −114. Nonpolar covalent: I₂ 114 · naphthalene 80 · paraffin wax about 60 · CO₂ sublimes at −78 · CH₄ −182 · N₂ −210 · O₂ −218.

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